Chemical Properties of Elements for HSC Chemistry

Suppose you’re given four unlabelled samples: sodium, magnesium, chlorine, and argon. They’re all elements, but putting each one into a reaction would give wildly different results. Sodium reacts violently with water. Magnesium reacts much more slowly. Chlorine readily forms compounds with metals. Argon mostly refuses to react at all.

The useful question isn’t just “what is each element?” It’s: can we predict that behaviour before doing the experiment?

Usually, yes. An element’s position in the periodic table tells us about its outer electrons, and those outer electrons strongly influence how it reacts.

Before reading on, predict this: which should be more reactive with water, lithium or potassium? They’re both in Group 1, but potassium is further down the group.

Potassium is more reactive. By the end of this guide, you should be able to explain why rather than simply memorising the trend.

01Chemical properties describe what an element does in a reaction

A chemical property describes how a substance behaves when it undergoes a chemical change.

For an element, useful chemical properties include:

• whether it reacts with water
• whether it reacts with oxygen
• whether it reacts with acids
• whether it tends to lose or gain electrons
• what ions it commonly forms
• how reactive it is compared with similar elements

This is different from a physical property, such as density, melting point, colour, or electrical conductivity. If you need to refresh that distinction, see Physical Properties of Elements for HSC Chemistry.

The periodic table helps with both, but chemical properties are especially connected to an atom’s valence electrons.

02Start with the outer electrons

An atom contains a positively charged nucleus surrounded by electrons.

The electrons in the highest occupied energy level are called valence electrons. These are the electrons most directly involved in chemical reactions.

You can picture an atom as having inner electrons tucked safely inside, while the valence electrons are standing near the front door. When another atom comes along, those outer electrons are the ones available for swapping, sharing, or leaving.

That picture is useful, but it is simplified. Electrons do not literally sit in circular shells like tiny planets. In chemistry, the shell model is a convenient way to understand patterns in electron energy and chemical behaviour.

For the main-group elements, the periodic table gives us an immediate clue about valence electrons.

This table is a pattern, not an absolute rule for every possible compound.

Check your understanding

Why are sodium and potassium chemically similar even though potassium has more electrons overall?

03Why atoms lose, gain, or share electrons

For many main-group elements, reactions can be understood using a simple idea: atoms often move towards a more stable outer-electron arrangement.

For example:

• sodium has one valence electron and can lose it
• chlorine has seven valence electrons and can gain one
• neon already has a filled outer shell, so it has little tendency to gain or lose electrons

So sodium and chlorine fit together neatly:

\[
\ce{Na -> Na+ + e-}
\]

\[
\ce{Cl + e- -> Cl-}
\]

The oppositely charged ions can then form an ionic compound, sodium chloride.

This is sometimes explained using the octet rule, where atoms tend towards eight electrons in their outer shell.

The octet rule is useful, especially for elements near the beginning of HSC chemistry, but don’t treat it as a universal law. There are important exceptions, particularly once you study more complicated molecules and elements beyond the first few periods.

04Metals and non-metals behave differently

One broad way to classify elements chemically is into metals, non-metals, and metalloids.

The distinction is not based on only one property. Instead, we look for patterns of behaviour.

Metals

Metals generally tend to lose electrons and form positive ions called cations.

For example:

\[
\ce{Mg -> Mg^2+ + 2e-}
\]

Many metals also react with acids to produce hydrogen gas.

For magnesium:

\[
\ce{Mg(s) + 2HCl(aq) -> MgCl2(aq) + H2(g)}
\]

Magnesium has been oxidised from neutral magnesium atoms to \(\mathrm{Mg}^{2+}\) ions.

More reactive metals generally lose electrons more easily.

Non-metals

Non-metals more commonly gain electrons or share electrons.

Chlorine, for example, can gain one electron:

\[
\ce{Cl + e- -> Cl-}
\]

Two non-metals can also share electrons to form covalent bonds. Oxygen atoms, for example, share electrons in an \(\ce{O2}\) molecule.

So don’t memorise “non-metals gain electrons” as though that is all they ever do. They frequently share electrons as well.

Metalloids

Metalloids sit near the boundary between metals and non-metals on the periodic table.

Elements commonly classified as metalloids include silicon and germanium. Their properties are intermediate in several ways, so they do not fit neatly into a simple metal/non-metal division.

For HSC questions, it is usually more useful to explain the relevant property than to rely only on the label “metalloid”.

05Group 1 metals: reactivity increases down the group

Now return to our original lithium versus potassium prediction.

Group 1 contains the alkali metals, including lithium, sodium, and potassium.

They all have one valence electron.

Their general reaction with water is:

\[
\ce{2M(s) + 2H2O(l) -> 2MOH(aq) + H2(g)}
\]

where \(M\) represents a Group 1 metal.

For sodium:

\[
\ce{2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g)}
\]

The key chemical change is that each sodium atom loses one electron to form \(\mathrm{Na}^{+}\).

Why does reactivity increase down Group 1?

Compare lithium and potassium.

Potassium has more occupied electron shells.

Its valence electron is:

1. further from the positive nucleus
2. shielded by more inner electrons
3. therefore less strongly attracted to the nucleus

So potassium loses its valence electron more easily than lithium.

That makes potassium more reactive.

The trend is:

\[
\ce{Li < Na < K < Rb < Cs}
\]

for Group 1 metal reactivity.

Application: predicting an unknown Group 1 metal

Suppose an unknown element is directly below sodium in Group 1.

Predict how its reaction with water would compare with sodium.

06Group 17 elements do the opposite

Group 17 contains the halogens, including fluorine, chlorine, bromine, and iodine.

A halogen atom has seven valence electrons. Instead of losing one electron, it tends to gain one electron.

For chlorine:

\[
\ce{Cl + e- -> Cl-}
\]

Now predict the trend.

If gaining an electron becomes harder as the outer shell gets further from the nucleus, would fluorine or iodine be more reactive?

Fluorine is more reactive.

This means the Group 17 trend runs in the opposite direction to Group 1:

\[
\ce{F > Cl > Br > I}
\]

in chemical reactivity.

Why does halogen reactivity decrease down the group?

As you move down Group 17:

• the atoms become larger
• there are more inner electron shells
• shielding increases
• an incoming electron is further from the nucleus
• attraction between the nucleus and that incoming electron becomes weaker

So the atom becomes less effective at gaining an electron.

This gives a useful contrast:

Don’t memorise these as two disconnected facts. The opposite trends happen because the two groups are trying to do opposite things.

Group 1 metals react by losing an electron.

Group 17 halogens react by gaining an electron.

07Application: halogen displacement

The reactivity trend lets us predict whether one halogen can displace another from a compound.

Consider chlorine gas added to potassium bromide solution:

\[
\ce{Cl2(g) + 2KBr(aq) -> 2KCl(aq) + Br2(aq)}
\]

Why does this happen?

Chlorine is above bromine in Group 17, so chlorine is more reactive. Chlorine gains electrons more readily than bromine.

The important ionic change is:

\[
\ce{Cl2 + 2Br- -> 2Cl- + Br2}
\]

Chlorine has displaced bromine.

Check your understanding

Would bromine displace chlorine from sodium chloride solution?

08Group 18: why noble gases barely react

Group 18 contains the noble gases, including helium, neon, and argon.

These elements have filled outer electron shells.

That makes them unusually stable compared with neighbouring elements.

They have little tendency to:

• lose electrons
• gain electrons
• form ordinary chemical bonds

So their chemical reactivity is very low.

This is why argon can be used where a relatively unreactive atmosphere is needed.

The simplified idea is “full outer shell equals stable”. That works well for explaining introductory periodic trends. It does not mean noble gases are literally incapable of forming compounds. Some heavier noble gases, particularly xenon, can form compounds under suitable conditions.

09Group 2 metals follow a similar pattern to Group 1

Group 2 elements have two valence electrons and commonly form \(2+\) ions.

For example:

\[
\ce{Mg -> Mg^2+ + 2e-}
\]

Like Group 1, their metallic reactivity generally increases as you move down the group because the valence electrons become easier to remove.

For example, calcium reacts more readily with water than magnesium under ordinary conditions.

But be careful with comparisons between different groups.

It is tempting to say:

“Anything further down the periodic table must be more reactive.”

That is wrong.

Position only becomes useful when you understand what reaction the element is undergoing.

Potassium becomes more reactive than lithium down Group 1 because losing an electron becomes easier.

Iodine becomes less reactive than chlorine down Group 17 because gaining an electron becomes harder.

There is no single “reactivity increases this way” arrow that works for the whole periodic table.

10What happens across a period?

Moving from left to right across a period, the number of protons increases while electrons are added to the same main energy level.

This generally causes the nucleus to attract the outer electrons more strongly.

As a broad pattern:

• metallic behaviour decreases
• non-metallic behaviour increases
• atoms become less willing to lose electrons
• towards the right side, atoms become more likely to attract electrons in bonding

Consider Period 3:

\[
\ce{Na,\ Mg,\ Al,\ Si,\ P,\ S,\ Cl,\ Ar}
\]

Sodium on the left is strongly metallic and readily forms \(\mathrm{Na}^{+}\).

Chlorine near the right is strongly non-metallic and readily forms \(\mathrm{Cl}^{-}\).

Argon at the far right is very unreactive because its outer shell is filled.

So a period can be thought of as a gradual change in the balance between losing electrons, sharing electrons, gaining electrons, and finally having little reason to react at all.

That is a useful mental model, although real chemistry becomes more complicated in the middle of the periodic table.

11A practical decision rule for HSC questions

If you’re asked to predict an element’s chemical behaviour from its periodic-table position, don’t jump straight to a memorised trend.

Use this sequence.

1. Identify the group.
   This gives you a first clue about the number of valence electrons.

2. Ask what the atom is likely to do with electrons.
   Does it tend to lose, gain, or share them?

3. If comparing elements in the same group, consider distance and shielding.
   More shells generally mean greater atomic radius and more shielding.

4. Connect that change to the actual reaction.
   Does the reaction require the atom to lose an electron or gain one?

5. Only then state the reactivity trend.

This prevents one of the most common mistakes: remembering the direction of a trend but forgetting why it points that way.

12Example: sodium versus magnesium

Predict which element is more likely to lose its first electron easily: sodium or magnesium.

Both are in Period 3.

Sodium is in Group 1 and has one valence electron. Magnesium is in Group 2 and has two.

Across the period from sodium to magnesium, nuclear charge increases. The outer electrons experience stronger attraction to the nucleus.

So sodium’s outer electron is easier to remove.

This helps explain why sodium is generally more reactive than magnesium in reactions that depend on losing electrons, such as reaction with cold water.

Sodium:

\[
\ce{2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g)}
\]

Magnesium reacts much less readily with cold water.

The useful conclusion is not simply “sodium is more reactive”. It is:

sodium more readily loses an electron because its valence electron is held less strongly.

That explanation can be transferred to unfamiliar questions.

13Example: chlorine versus iodine

An HSC question might tell you that chlorine and iodine both react by gaining electrons and then ask which is more reactive.

Both are Group 17 elements, but iodine is further down the group.

Iodine has:

• more occupied electron shells
• a larger atomic radius
• greater shielding

Its nucleus therefore attracts an incoming electron less strongly.

So chlorine gains an electron more readily and is more reactive than iodine.

This predicts the reaction:

\[
\ce{Cl2 + 2I- -> 2Cl- + I2}
\]

but not the reverse reaction under the same conditions.

14The misconception to avoid: “reactivity means wanting a full shell”

Students often say something like:

“Potassium is reactive because it really wants a full outer shell.”

That points in roughly the right direction, but it is not a strong chemical explanation.

Atoms don’t “want” anything.

A better explanation describes the actual energetic and electrostatic situation.

For potassium, the outer electron is relatively far from the nucleus and strongly shielded by inner electrons. The attraction between the nucleus and that electron is therefore weaker, so the electron is more easily removed.

Likewise, saying “chlorine wants one electron” is shorthand.

The more precise idea is that chlorine can attract an additional electron strongly enough for electron gain to be favourable in many reactions.

Use the simple language to build intuition, then upgrade it when you explain.

15A final prediction

An unknown element \(X\):

• is in Group 2
• is below magnesium
• forms \(X^{2+}\) ions
• reacts with water more readily than magnesium

Does that information fit the periodic trend?

Yes.

A Group 2 element below magnesium has more occupied electron shells. Its valence electrons are further from the nucleus and more shielded. They are therefore easier to remove, so increased metallic reactivity is expected.

If \(X\) is immediately below magnesium, it is calcium.

You can now go one step further than simply classifying an element as “metal” or “non-metal”. Periodic-table position lets you predict what kind of ion it is likely to form, whether it tends to lose or gain electrons, and how its reactivity should change compared with neighbouring elements.

That electron behaviour is the bridge to the next major idea: predicting the formulas and names of compounds formed when elements react. The same group patterns used here become especially useful when working with IUPAC Naming of Inorganic Substances for HSC Chemistry.