Atomic Properties Change Predictably Across Periods and Down Groups
- 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.
- Three key properties you should know are:
- Atomic radius
- Ionization energy
- Electronegativity
- These trends help explain why some elements are more reactive than others.
Atomic Radius
- Atomic radius is a measure of the size of an atom.
- It is often defined as half the distance between the nuclei of two bonded atoms of the same element.
Analogy
You can also think of it more simply as the distance from the nucleus to the outermost electrons.
Atomic Radius Increases Down a Group
As you move down a group in the periodic table:
- Atoms gain more electron shells, so the distance between the nucleus and the outer electrons increases.
- The nuclear charge (number of protons) increases, but this effect is partly cancelled by extra inner electron shells, which provide shielding.
- Because of greater shielding and larger distance, the outer electrons are held less tightly.
Therefore, atomic radius increases down a group.
Example
In Group 1:
- Lithium (Li) has a smaller atomic radius than cesium (Cs).
- Cesium has more electron shells, so its atoms are much larger.
Atomic Radius Decreases Across a Period
As you move across a period from left to right:
- The number of protons in the nucleus increases → stronger nuclear charge.
- Electrons are added to the same main energy level (same shell).
- 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
Definition
Ionisation energy
Ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms.
Note
For MYP, you can think of it as how much energy it takes to remove an outer electron from an atom.
Ionization Energy Decreases Down a Group
As you move down a group:
- Outer electrons are farther from the nucleus.
- There are more inner shells, so shielding increases.
- 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.
Example
As you move down a group:
- Outer electrons are farther from the nucleus.
- There are more inner shells, so shielding increases.
- 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:
- The nuclear charge increases (more protons).
- Electrons are added to the same shell, so shielding remains similar.
- 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.
Example
In Period 2:
- Fluorine (F) has a higher ionisation energy than lithium (Li).
- Fluorine has a greater nuclear charge and a smaller atomic radius, so its electrons are held more tightly.
Electronegativity
Definition
Electronegativity
The ability of an atom to attract a bonding pair of electrons in a covalent bond.
Hint
It is a relative scale (we compare how strongly different atoms attract electrons when they are bonded).
Electronegativity Decreases Down a Group
As you move down a group:
- Atoms get larger (greater atomic radius).
- The bonding electrons are, on average, further from the nucleus.
- Increased shielding reduces the effective pull of the nucleus on bonding electrons.
So electronegativity decreases down a group.
Example
In Group 17 (halogens):
- Chlorine (Cl) is more electronegative than bromine (Br).
- Chlorine has a smaller atomic radius, so it attracts bonding electrons more strongly.
Electronegativity Increases Across a Period
As you move across a period from left to right:
- The nuclear charge increases (more protons).
- Electrons are added to the same shell, so shielding does not increase much.
- Atoms become smaller, and the nucleus attracts bonding electrons more strongly.
Therefore, electronegativity increases across a period.
Example
In Period 2:
- Oxygen (O) is more electronegative than carbon (C).
- Oxygen has a higher nuclear charge and a similar shell structure, so it attracts bonding electrons more strongly.
Melting Point
- Melting and boiling points are closely linked to the type of bonding and the strength of forces between particles (atoms, ions or molecules).
- Across the periodic table, we can see clear trends.
Melting Points Decrease Down Group 1
- For the Group 1 metals (Li, Na, K, Rb, Cs), the melting points decrease as you go down the group.
- Group 1 metals are held together by metallic bonding: positive metal ions in a lattice are attracted to a “sea” of delocalised electrons.
- As you go down the group:
- The metal ions become larger (more electron shells).
- The charge stays the same (+1), so the charge density (charge per unit size) decreases.
- The attraction between the positive ions and the delocalised electrons becomes weaker.
- Weaker metallic bonding means less energy is needed to break the lattice and melt the metal, so the melting point falls.
Example
- Lithium (Li) has a higher melting point and is harder than cesium (Cs).
- Cesium is very soft and has a low melting point; it can almost melt in your hand.
Melting Points Increase Down Group 17
- For the Group 17 elements (the halogens: F₂, Cl₂, Br₂, I₂, At₂):
- The melting and boiling points increase as you go down the group.
- The halogens also become less volatile (harder to turn into a gas).
- Halogens exist as simple covalent molecules (F₂, Cl₂, Br₂, I₂).
- Between the molecules there are intermolecular forces, mainly London dispersion forces (van der Waals forces).
- As you go down the group:
- The molecules get larger (more electrons).
- The electron cloud becomes more easily polarised.
- This increases the strength of the intermolecular forces.
- 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.
- On the left: metals (e.g. Na, Mg, Al)
- Have metallic bonding.
- Generally have higher melting points, because many positive ions are strongly attracted to a sea of delocalized electrons.
- In the middle (for some periods): giant covalent structures (e.g. Si in Period 3)
- Each atom is joined by strong covalent bonds in a giant network.
- Very high melting points (lots of energy needed to break many strong covalent bonds).
- On the right: simple molecular substances and noble gases (e.g. P₄, S₈, Cl₂, Ar)
- Particles are molecules or single atoms held together by weak intermolecular forces.
- Low melting and boiling points.
Hint
So across a period, there is a trend from:
Metallic structures → giant covalent network → simple molecules / individual atoms
and the melting point usually:
Rises to a maximum in the giant covalent structure (like Si), then falls sharply for the simple molecular substances.
Example
In Period 3:
- Sodium (Na) is a metal with a relatively high melting point (metallic bonding).
- Silicon (Si) has an even higher melting point (giant covalent structure).
- Chlorine (Cl₂) is a simple molecular substance with low melting point, and argon (Ar), a noble gas, has a very low melting point.
Reactivity of Group 1 Elements Increases Down the Group
- The reactivity of Group 1 metals increases as you go down the group.
- 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⁺).
- As you go down Group 1:
- Atoms become larger (more shells, larger atomic radius).
- The single outer electron is further from the nucleus.
- There are more inner electron shells, so shielding increases.
- The attraction between the nucleus and the outer electron becomes weaker.
- Therefore, the outer electron is more easily removed, so reactivity increases down the group.
Example
- Potassium (K) reacts more vigorously with water than lithium (Li).
- Lithium fizzes gently; potassium ignites the hydrogen, often producing a lilac flame.
Reactivity of Group 17 Elements Decreases Down the Group.
- The reactivity of Group 17 (halogens) decreases as you go down the group.
- Halogens gain one electron in reactions to form X⁻ ions (e.g. Cl⁻, Br⁻, I⁻).
- As you go down Group 17:
- Halogen atoms become larger (greater atomic radius).
- The outer shell is further from the nucleus.
- There is more shielding from inner shells.
- The attraction between the nucleus and an incoming electron becomes weaker.
- So it is harder for the atom to gain an electron, and reactivity decreases down the group.
Example
- Chlorine (Cl₂) is a green gas at room temperature.
- Bromine (Br₂) is a reddish-brown liquid.
- Iodine (I₂) is a purple-black solid.
- This change in physical state reflects increasing melting and boiling points and stronger intermolecular forces down the group.
Example
- Chlorine reacts vigorously with sodium to form sodium chloride (NaCl), common table salt, important in food preservation and flavouring.
- Chlorine is also widely used in water purification, helping to kill harmful microorganisms and provide safe drinking water.
Note
Everyday uses of halogens:
- Chlorine: Disinfecting drinking water, sanitising swimming pools, household bleach.
- Fluorine (as fluoride ions): In toothpaste to reduce tooth decay.
- Bromine: Used in certain flame retardants, photographic chemicals and medicines.
- Iodine: Added to table salt to prevent iodine deficiency, and used as an antiseptic on wounds.
Group 18 Elements Are Inert and Nonreactive
- Group 18 elements (the noble gases) are very stable because they have full valence shells.
- They do not need to gain, lose or share electrons.
- They are therefore chemically inert, and rarely form compounds.
Example
Helium (He) is used to fill balloons because it is very unreactive and non-flammable.
Exam technique
Quick Summary of Trends
- Atomic radius
- Increases down a group
- Decreases across a period
- Ionisation energy
- Decreases down a group
- Increases across a period
- Electronegativity
- Decreases down a group
- Increases across a period
- Melting / boiling points
- Group 1 (alkali metals):
- Melting points decrease down the group (metallic bonds weaken as ions get larger).
- Group 17 (halogens):
- Melting and boiling points increase down the group (larger molecules → stronger intermolecular forces).
- Across a period:
- Generally higher for metals and giant covalent structures, then drop sharply for simple molecular substances and noble gases.
- Group 1 (alkali metals):
- Reactivity
- Group 1 (alkali metals):
- Reactivity increases down the group (outer electron further from nucleus, more shielding, easier to remove).
- Group 17 (halogens):
- Reactivity decreases down the group (harder to gain an electron as atoms get larger and more shielded).
- Group 18 (noble gases):
- Very low reactivity (full valence shells, very stable).
- Group 1 (alkali metals):
Self review
- Why does atomic radius increase down a group but decrease across a period
- Explain why rubidium has a lower ionization energy than sodium.
- Why is oxygen more electronegative than carbon?
- How do these trends help explain why Group 1 metals are very reactive and Group 17 halogens are also very reactive, but in a different way?
- Why do melting points decrease down Group 1, but increase down Group 17
- Explain, in terms of electron shells and shielding, why Group 1 metals become more reactive down the group.
- Explain why Group 17 halogens become less reactive down the group.
- How does the type of bonding explain the trend in melting points across a period?
