MYP Integrated Sciences Periodic Trends Notes

Noble Gases

1. The noble gases are the elements in Group 18 of the periodic table (sometimes called Group 0 in older notation).
2. They include helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn).
3. They are often described as “inert” or unreactive because they rarely form compounds.
4. To understand why, we need to look at their electron configurations.

Why Noble Gases Are Unreactive?

1. Noble gases have full outer (valence) electron shells.
2. A full valence shell is a very stable arrangement of electrons.
3. Because their outer shells are already full, noble gases do not need to gain, lose or share electrons.
4. As a result, they very rarely react with other elements.

These configurations all show completely filled outer shells, which explains why these atoms are so unreactive under normal conditions.

Properties of Noble Gases

The unique electron configurations of noble gases lead to several important properties:

1. Very low reactivity:
   1. As discussed earlier, noble gases have full valence shells, which makes them very stable.
   2. They do not need to gain, lose or share electrons, so they rarely form compounds.
2. Monatomic and non-polar:
   1. Noble gases exist as single atoms (monatomic gases), such as He(g), Ne(g) and Ar(g).
   2. Because they are not bonded to other atoms, they are non-polar and interact only weakly with each other.
3. Low melting and boiling points:
   1. There are only weak van der Waals (dispersion) forces between noble gas atoms.
   2. As a result, they have low melting and boiling points and are gases at room temperature.

Uses of Noble Gases

Because of their inertness and physical properties, noble gases are very useful in many practical applications.

Lighting

1. So-called “neon lights” can contain different noble gases, not just neon.
2. When an electric current passes through the gas at low pressure, the atoms are excited and then emit light as they return to lower energy levels.
3. Different noble gases produce different colours.

Welding

1. Argon is commonly used as a shielding gas in welding.
2. Its inert nature prevents the hot metal from reacting with oxygen or nitrogen in the air.
3. This protects the weld from contamination and improves the quality of the joint.

Preserving Food

1. Noble gases such as argon and helium can be used to displace oxygen in food packaging.
2. Less oxygen means less oxidation and slower growth of many microorganisms.
3. This helps to extend the shelf life of perishable foods.
4. In industry, nitrogen is also commonly used for this purpose.

The Importance of Understanding Noble Gases in the Atmosphere

Even though are noble gases unreactive, monitoring their presence in the atmosphere is important.

Argon in the Atmosphere

1. Argon (Ar) makes up about 0.93% of Earth’s atmosphere.
2. It is the third most abundant gas, after nitrogen (N₂) and oxygen (O₂).
3. Scientists use argon isotopes (for example in K–Ar dating) to estimate the age of rocks and minerals.

Helium Shortage

1. Helium (He) is used in many applications:
   1. Filling balloons and airships (light and non-flammable)
   2. Cooling superconducting magnets in MRI scanners and other scientific instruments
2. Helium is produced in limited amounts (e.g. from radioactive decay in rocks and extracted from natural gas) and can be lost to space.
3. Therefore, helium is a valuable resource that needs careful management.

Noble Gases as Tracers

Because noble gases are chemically inert, they can act as excellent tracers:

1. They can be used to track the movement of air masses in the atmosphere.
2. In ice cores and groundwater, dissolved noble gases help scientists estimate the age of the samples and the conditions under which they formed.
3. This information contributes to our understanding of climate change and helps in predicting future environmental trends.

Atomic Properties Change Predictably Across Periods and Down Groups

1. Many atomic properties change in regular patterns as you move across a period (left to right) or down a group (top to bottom) in the periodic table.
2. Three key properties you should know are:
   1. Atomic radius
   2. Ionization energy
   3. Electronegativity
3. These trends help explain why some elements are more reactive than others.

Atomic Radius

1. Atomic radius is a measure of the size of an atom.
2. It is often defined as half the distance between the nuclei of two bonded atoms of the same element.

Atomic Radius Increases Down a Group

As you move down a group in the periodic table:

1. Atoms gain more electron shells, so the distance between the nucleus and the outer electrons increases.
2. The nuclear charge (number of protons) increases, but this effect is partly cancelled by extra inner electron shells, which provide shielding.
3. Because of greater shielding and larger distance, the outer electrons are held less tightly.

Therefore, atomic radius increases down a group.

Atomic Radius Decreases Across a Period

As you move across a period from left to right:

1. The number of protons in the nucleus increases → stronger nuclear charge.
2. Electrons are added to the same main energy level (same shell).
3. There is little or no increase in shielding from inner shells.

The stronger nuclear attraction pulls the outer electrons closer to the nucleus, so the atomic radius decreases across a period.

Ionization Energy

Ionization Energy Decreases Down a Group

As you move down a group:

1. Outer electrons are farther from the nucleus.
2. There are more inner shells, so shielding increases.
3. The attraction between the nucleus and the outer electrons is weaker.

This means it is easier to remove an electron, so ionisation energy decreases down a group.

Ionization Energy Increases Across a Period

As you move across a period from left to right:

1. The nuclear charge increases (more protons).
2. Electrons are added to the same shell, so shielding remains similar.
3. The attraction between the nucleus and the outer electrons becomes stronger.

Therefore, it becomes harder to remove an electron, and ionisation energy increases across a period.

Electron Affinity

Electron Affinity Becomes More Negative Across a Period

As you move across a period from left to right:

1. The nuclear charge increases (more protons).
2. Electrons are added to the same shell, so shielding changes only a little.
3. Atoms get smaller, so the added electron is closer to the nucleus.
4. The attraction between the nucleus and the added electron becomes stronger.

This means more energy is released when the electron is gained, so electron affinity generally becomes more negative across a period.

Electron Affinity Becomes Less Negative Down a Group

As you move down a group:

1. The outer shell is farther from the nucleus.
2. There are more inner shells, so shielding increases.
3. The added electron feels less attraction to the nucleus.

This means less energy is released when an electron is gained, so electron affinity generally becomes less negative down a group.

Electronegativity

Electronegativity Decreases Down a Group

As you move down a group:

1. Atoms get larger (greater atomic radius).
2. The bonding electrons are, on average, further from the nucleus.
3. Increased shielding reduces the effective pull of the nucleus on bonding electrons.

So electronegativity decreases down a group.

Electronegativity Increases Across a Period

As you move across a period from left to right:

1. The nuclear charge increases (more protons).
2. Electrons are added to the same shell, so shielding does not increase much.
3. Atoms become smaller, and the nucleus attracts bonding electrons more strongly.

Therefore, electronegativity increases across a period.

Melting Point

1. Melting and boiling points are closely linked to the type of bonding and the strength of forces between particles (atoms, ions or molecules).
2. Across the periodic table, we can see clear trends.

Melting Points Decrease Down Group 1

1. For the Group 1 metals (Li, Na, K, Rb, Cs), the melting points decrease as you go down the group.
2. Group 1 metals are held together by metallic bonding: positive metal ions in a lattice are attracted to a “sea” of delocalised electrons.
3. As you go down the group:
   1. The metal ions become larger (more electron shells).
   2. The charge stays the same (+1), so the charge density (charge per unit size) decreases.
   3. The attraction between the positive ions and the delocalised electrons becomes weaker.
4. Weaker metallic bonding means less energy is needed to break the lattice and melt the metal, so the melting point falls.

Melting Points Increase Down Group 17

1. For the Group 17 elements (the halogens: F₂, Cl₂, Br₂, I₂, At₂):
2. The melting and boiling points increase as you go down the group.
3. The halogens also become less volatile (harder to turn into a gas).
4. Halogens exist as simple covalent molecules (F₂, Cl₂, Br₂, I₂).
5. Between the molecules there are intermolecular forces, mainly London dispersion forces (van der Waals forces).
6. As you go down the group:
   1. The molecules get larger (more electrons).
   2. The electron cloud becomes more easily polarised.
   3. This increases the strength of the intermolecular forces.
7. Stronger intermolecular forces mean more energy is needed to separate the molecules, so melting and boiling points rise down the group.

Melting Points Vary Across a Period

The pattern of melting points across a period is linked to changes in bonding and structure.

1. On the left: metals (e.g. Na, Mg, Al)
   1. Have metallic bonding.
   2. Generally have higher melting points, because many positive ions are strongly attracted to a sea of delocalized electrons.
2. In the middle (for some periods): giant covalent structures (e.g. Si in Period 3)
   1. Each atom is joined by strong covalent bonds in a giant network.
   2. Very high melting points (lots of energy needed to break many strong covalent bonds).
3. On the right: simple molecular substances and noble gases (e.g. P₄, S₈, Cl₂, Ar)
   1. Particles are molecules or single atoms held together by weak intermolecular forces.
   2. Low melting and boiling points.

Reactivity of Group 1 Elements Increases Down the Group

1. The reactivity of Group 1 metals increases as you go down the group.
2. Group 1 metals all have one electron in their outer shell. In reactions, they lose this electron to form M⁺ ions (e.g. Na⁺, K⁺).
3. As you go down Group 1:
   1. Atoms become larger (more shells, larger atomic radius).
   2. The single outer electron is further from the nucleus.
   3. There are more inner electron shells, so shielding increases.
   4. The attraction between the nucleus and the outer electron becomes weaker.
4. Therefore, the outer electron is more easily removed, so reactivity increases down the group.

Reactivity of Group 17 Elements Decreases Down the Group.

1. The reactivity of Group 17 (halogens) decreases as you go down the group.
2. Halogens gain one electron in reactions to form X⁻ ions (e.g. Cl⁻, Br⁻, I⁻).
3. As you go down Group 17:
   1. Halogen atoms become larger (greater atomic radius).
   2. The outer shell is further from the nucleus.
   3. There is more shielding from inner shells.
   4. The attraction between the nucleus and an incoming electron becomes weaker.
4. So it is harder for the atom to gain an electron, and reactivity decreases down the group.

Group 18 Elements Are Inert and Nonreactive

1. Group 18 elements (the noble gases) are very stable because they have full valence shells.
2. They do not need to gain, lose or share electrons.
3. They are therefore chemically inert, and rarely form compounds.