Using the Periodic Table
3. Predicting Ion Formation
Learning outcomes
- I can use periodic trends to predict how elements will react.
- I can compare the expected behavior of elements from different groups.
- I can explain how electron arrangements influence chemical reactivity.
- I can use periodic table information to predict reaction outcomes.
- I can justify predictions using evidence from periodic trends.
Using the Periodic Table to Predict Reactions
The periodic table is a powerful predictive model.
An element's position gives us information about its:
- valence electrons
- tendency to gain or lose electrons
- likely ion charge
- atomic radius
- ionization energy
- electronegativity
- chemical reactivity
By combining these patterns, we can make predictions about elements that we may never have studied directly.
The central reasoning chain is:
periodic-table position → electron arrangement → electron gain or loss → ion formation → reactivity → predicted reaction
What Happens When an Ion Forms?
An ion forms when an atom gains or loses electrons.
If an atom loses electrons:
positive ion → cation
If an atom gains electrons:
negative ion → anion
The number of protons does not change during ordinary chemical ion formation.
Only the number of electrons changes.
Why Electron Arrangement Matters
The electrons most important in chemical reactions are the valence electrons.
These are the electrons in the outermost occupied shell.
Many main-group atoms react in ways that produce a more stable outer electron arrangement.
For many introductory examples, this means reaching an electron arrangement similar to a nearby noble gas.
Metals Usually Lose Electrons
Metals are mainly found on the left and centre of the periodic table.
Many main-group metals have relatively few valence electrons.
Examples:
Sodium:
2,8,1
Magnesium:
2,8,2
Aluminium:
2,8,3
These atoms can reach stable outer-shell arrangements by losing electrons.
Therefore:
metals generally form positive ions.
Non-Metals Often Gain Electrons
Reactive non-metals toward the right side of the periodic table often have nearly complete outer shells.
Examples:
Nitrogen:
2,5
Oxygen:
2,6
Chlorine:
2,8,7
These elements can reach complete outer shells by gaining electrons.
Therefore:
reactive non-metals often form negative ions.
Noble Gases Are Different
Noble gases already have complete outer electron shells.
Neon:
2,8
Argon:
2,8,8
Because these arrangements are already stable, noble gases have little tendency to gain or lose electrons.
Therefore, they are generally:
very unreactive
and do not commonly form simple ions.
Predicting Common Ion Charges
Group membership provides a useful pattern for many main-group elements.
Group 1 → +1
Group 2 → +2
Group 13 → +3
Group 15 → −3
Group 16 → −2
Group 17 → −1
Group 18 → generally no common simple ion
These charges tell us what ions are likely to form, but periodic trends help us explain how readily those ions form.
Periodic Trends and Reactivity
Several periodic trends are especially useful when predicting chemical behaviour:
- atomic radius
- shielding
- ionization energy
- electronegativity
These properties are connected.
Understanding these trends is more useful than simply memorizing which element is "more reactive."
Atomic Radius
Atomic radius describes the size of an atom.
Down a group:
atomic radius generally increases.
Why?
Each new period adds another occupied electron shell.
The outer electrons therefore become farther from the nucleus.
This can significantly affect chemical reactivity.
Electron Shielding
Inner electrons reduce the effective attraction between the nucleus and outer electrons.
This effect is called shielding.
Down a group:
- more occupied shells are present
- more inner electrons are present
- shielding generally increases
As a result, outer electrons can experience less effective attraction from the nucleus.
Ionization Energy
Ionization energy is the energy required to remove an electron from an isolated gaseous atom.
For metals, this is especially important because metals react by losing electrons.
Lower ionization energy means:
electron removal is easier.
Higher ionization energy means:
electron removal is harder.
Therefore, ionization energy can help us predict how readily a metal forms positive ions.
Electronegativity
Electronegativity describes an atom's tendency to attract bonding electrons.
Electronegativity generally:
increases across a period from left to right
and:
decreases down a group
Highly electronegative non-metals strongly attract electrons in chemical bonds.
This helps explain the behaviour of elements such as fluorine and chlorine.
Group 1: Alkali Metals
Group 1 elements have:
one valence electron
Examples:
Li → 2,1
Na → 2,8,1
K → 2,8,8,1
They tend to lose this electron:
M → M⁺ + e⁻
Therefore, Group 1 metals form:
+1 ions
Group 1 Reactivity Trend
Moving down Group 1:
Li → Na → K → Rb → Cs
reactivity generally:
increases
Why?
Down the group:
- atomic radius increases
- shielding increases
- the outer electron is farther from the nucleus
- first ionization energy generally decreases
- the outer electron becomes easier to remove
Therefore:
+1 ions form more readily during reactions.
Predicting Group 1 Reactions
Suppose an unfamiliar element X is below potassium in Group 1.
We can predict that X:
- has one valence electron
- forms X⁺
- loses an electron during reactions
- is highly reactive
- should generally be more reactive than potassium
This prediction comes from group position and periodic trends, not from knowing the identity of X.
Group 1 and Water
Group 1 metals react with water.
General equation:
2M + 2H₂O → 2MOH + H₂
For example:
2Na + 2H₂O → 2NaOH + H₂
The reactions generally become more vigorous down the group.
Therefore:
potassium reacts more vigorously than sodium
and:
sodium reacts more vigorously than lithium.
Explaining the Group 1 Prediction
A strong explanation would be:
Potassium has more occupied electron shells than sodium.
Therefore, potassium has:
- greater atomic radius
- greater shielding
Its outer electron experiences weaker effective attraction to the nucleus.
It is therefore easier to remove.
So potassium forms K⁺ more readily during reactions and is more reactive than sodium.
This is much stronger than simply writing:
"Potassium is lower in the group."
Group 2 Metals
Group 2 metals have:
two valence electrons
They usually lose both:
M → M²⁺ + 2e⁻
Examples include:
Mg²⁺
Ca²⁺
Sr²⁺
Ba²⁺
Like Group 1, their chemical behaviour can often be understood in terms of electron loss.
Predicting Group 2 Behaviour
Suppose X is an unfamiliar Group 2 element.
We can predict that it:
- is metallic
- has two valence electrons
- tends to lose two electrons
- forms X²⁺
- can react with non-metals to form ionic compounds
If X is lower in the group, increased atomic radius and shielding generally make its outer electrons easier to remove.
Group 17: Halogens
Halogens have:
seven valence electrons
Examples:
F → 2,7
Cl → 2,8,7
They tend to gain one electron:
X + e⁻ → X⁻
Therefore, halogens commonly form:
−1 ions
Group 17 Reactivity Trend
Moving down Group 17:
F → Cl → Br → I
reactivity generally:
decreases
Why?
Down the group:
- atomic radius increases
- shielding increases
- the outer region is farther from the nucleus
- attraction for an incoming electron becomes less effective
Therefore, gaining an electron becomes less favourable.
Group 1 vs Group 17
These groups provide an important comparison.
Group 1
Reacts by:
losing an electron
Down the group, electron loss becomes easier.
Therefore:
reactivity increases.
Group 17
Reacts by:
gaining an electron
Down the group, attraction for the incoming electron becomes weaker.
Therefore:
reactivity decreases.
The same changes in radius and shielding produce opposite reactivity trends because the groups react in different ways.
Predicting Halogen Displacement Reactions
A more reactive halogen can displace a less reactive halogen from a compound containing its halide ions.
For the commonly studied halogens:
Cl > Br > I
in reactivity.
Therefore:
Cl₂ + 2KBr → 2KCl + Br₂
Chlorine displaces bromine.
But:
Br₂ + KCl → no reaction
because bromine is less reactive than chlorine.
Predicting an Unfamiliar Displacement
Suppose halogen X displaces bromide ions but does not displace chloride ions.
What can we conclude?
X must be:
more reactive than bromine
but:
less reactive than chlorine
Therefore:
Cl > X > Br
We can use reaction evidence to determine the relative position of an unfamiliar element in a reactivity series.
Noble Gases
Group 18 elements already have complete outer shells.
They therefore have little tendency to:
- lose electrons
- gain electrons
- form ordinary chemical bonds
Therefore, if an unfamiliar element is identified as a noble gas, our first prediction is:
very low chemical reactivity.
Transition Metals
Transition metals behave differently from simple Group 1 and Group 17 patterns.
They are metals and therefore generally form positive ions, but many can form more than one ion.
Examples:
Fe²⁺ and Fe³⁺
Cu⁺ and Cu²⁺
Their reactivity also varies considerably.
For example:
- iron reacts more readily than copper in many situations
- copper is relatively unreactive
- gold and platinum are especially resistant to many reactions
Therefore, transition metals cannot be described using one simple reactivity trend.
Comparing Different Groups
Consider sodium, magnesium, chlorine, and argon.
Sodium
Group 1.
One valence electron.
Tends to lose one.
Forms:
Na⁺
Highly reactive metal.
Magnesium
Group 2.
Two valence electrons.
Tends to lose two.
Forms:
Mg²⁺
Reactive metal, but behaviour differs from sodium.
Chlorine
Group 17.
Seven valence electrons.
Tends to gain one.
Forms:
Cl⁻
Reactive non-metal.
Argon
Group 18.
Complete outer shell.
Little tendency to gain or lose electrons.
Very unreactive.
Metals Reacting with Non-Metals
When a reactive metal reacts with a reactive non-metal, electron transfer often occurs.
For example:
Sodium loses an electron:
Na → Na⁺ + e⁻
Chlorine gains an electron:
Cl + e⁻ → Cl⁻
The resulting ions attract each other.
The product is:
NaCl
We can therefore use periodic-table position to predict both the ions and the compound.
Predicting Magnesium and Chlorine
Magnesium is in Group 2.
Predict:
Mg²⁺
Chlorine is in Group 17.
Predict:
Cl⁻
To balance charges:
one Mg²⁺ requires two Cl⁻.
Therefore:
MgCl₂
We can predict this formula without memorizing it.
Predicting Potassium and Oxygen
Potassium is Group 1.
Therefore:
K⁺
Oxygen is Group 16.
Therefore:
O²⁻
Two K⁺ ions are needed to balance one O²⁻:
2(+1) + (−2) = 0
Therefore:
K₂O
Predicting Calcium and Nitrogen
Calcium:
Group 2 → Ca²⁺
Nitrogen:
Group 15 → N³⁻
To balance:
3 Ca²⁺ → +6
2 N³⁻ → −6
Therefore:
Ca₃N₂
This demonstrates how periodic-table information can predict a reaction product's composition.
Reaction Outcomes and the Reactivity Series
Periodic position is extremely useful, but sometimes we also need a more specific reactivity series.
For example, one metal can sometimes displace another metal from a compound.
General pattern:
more reactive metal + less reactive metal compound → displacement
For example:
Zn + CuSO₄ → ZnSO₄ + Cu
Zinc is more readily oxidized than copper under these conditions.
Therefore, zinc displaces copper from the compound.
Using Experimental Evidence
Sometimes we can determine reactivity from experimental results rather than simply reading the periodic table.
Suppose:
Metal X displaces metal Y from a compound.
This tells us:
X is more reactive than Y
for the reaction being considered.
If:
X does not displace Z,
then Z may be more reactive than X.
We can combine these observations to build a reactivity order.
Example: Building a Reactivity Order
Suppose:
X displaces Y.
Y displaces Z.
Then:
X > Y > Z
in the relevant reactivity series.
If Z cannot displace X or Y, this supports the same order.
This type of reasoning allows us to solve unfamiliar problems using evidence rather than memorization.
Predicting Reaction or No Reaction
Suppose chlorine is added to potassium iodide.
Chlorine is more reactive than iodine.
Therefore, chlorine displaces iodide:
Cl₂ + 2KI → 2KCl + I₂
Prediction:
reaction occurs
Now suppose iodine is added to potassium chloride.
Iodine is less reactive than chlorine.
Prediction:
no displacement reaction
Using Atomic Structure to Justify Predictions
A good chemical prediction should include a reason.
Weak answer:
"Potassium is more reactive than lithium."
Better answer:
"Potassium is lower in Group 1."
Strong answer:
"Potassium has more occupied electron shells than lithium, giving it a larger atomic radius and greater shielding. Its outer electron experiences weaker effective nuclear attraction and is easier to remove, so potassium forms K⁺ more readily and is more reactive."
The strongest explanation connects the observation to atomic structure.
Predicting from Electron Arrangement
Suppose an unfamiliar element has:
2,8,1
We can immediately predict:
- one valence electron
- likely Group 1
- metallic behaviour
- electron loss
- +1 ion formation
- likely ionic bonding with non-metals
If the element reacts with chlorine, we would predict a compound with the general formula:
XCl
Another Electron-Arrangement Problem
An unfamiliar element has:
2,8,6
It has six valence electrons.
It is likely a Group 16 non-metal.
It needs two electrons to complete its outer shell.
Therefore:
X + 2e⁻ → X²⁻
If it reacts with a Group 1 metal M⁺, the predicted formula is:
M₂X
Comparing Two Unfamiliar Metals
Element X and element Y are both in Group 1.
X is below Y.
Which is likely to react more vigorously?
Prediction:
X
Reason:
X has more electron shells and greater shielding.
Its outer electron is easier to remove.
Therefore, X should form X⁺ more readily.
Comparing Two Unfamiliar Halogens
Elements A and B are both in Group 17.
A is above B.
Which is more reactive?
Prediction:
A
Reason:
A has fewer occupied shells, a smaller atomic radius, and less shielding.
Its nucleus attracts an incoming electron more effectively.
Therefore, A gains an electron more readily.
Predicting Metal + Halogen Reactions
Suppose an unfamiliar Group 2 metal X reacts with an unfamiliar Group 17 halogen Y.
Predict the ions:
X²⁺
Y⁻
Two Y⁻ ions are needed for every X²⁺ ion.
Therefore, the expected compound is:
XY₂
This prediction comes entirely from periodic-table information.
Predicting Group 1 + Group 16
Suppose X is a Group 1 metal.
Y is a Group 16 non-metal.
X forms:
X⁺
Y forms:
Y²⁻
Two X⁺ ions balance one Y²⁻.
Therefore:
X₂Y
Again, we can predict the product without knowing the actual elements.
Periodic Trends Across a Period
Across a period from left to right, several broad changes occur:
- atomic radius generally decreases
- first ionization energy generally increases
- electronegativity generally increases
- metallic character generally decreases
These changes help explain the transition from:
electron-losing metals
on the left toward:
electron-attracting non-metals
on the right.
From Metals to Non-Metals Across a Period
Consider Period 3:
Na → Mg → Al → Si → P → S → Cl → Ar
Toward the left:
elements tend to show stronger metallic behaviour and form positive ions.
Toward the right:
elements become increasingly non-metallic and tend to attract electrons more strongly.
At the far right:
argon has a complete outer shell and is very unreactive.
This pattern demonstrates periodicity.
Periodic Trends Are Evidence
When asked to justify a prediction, useful evidence includes:
Group membership
Tells us about valence electrons.
Atomic radius
Helps describe distance between the nucleus and outer electrons.
Shielding
Helps explain effective nuclear attraction.
Ionization energy
Helps predict how easily electrons can be removed.
Electronegativity
Helps predict how strongly atoms attract bonding electrons.
Known group trends
Allow comparisons between related elements.
A Prediction Framework
When solving an unfamiliar reaction problem, use this process.
Step 1: Locate the element.
Identify its group and period.
Step 2: Determine the valence-electron pattern.
How many outer electrons does it have?
Step 3: Decide whether it tends to gain or lose electrons.
Metal → usually lose.
Non-metal → often gain.
Step 4: Predict the ion.
Determine the likely charge.
Step 5: Consider periodic trends.
Think about:
- radius
- shielding
- ionization energy
- electronegativity
Step 6: Compare the elements.
Which should gain or lose electrons more readily?
Step 7: Predict the reaction outcome.
Will a reaction or displacement occur?
What ions or compounds are likely?
Step 8: Justify the prediction.
Connect the prediction to atomic structure and periodic evidence.
Worked Example 1
Which is more reactive: lithium or potassium?
Both are Group 1 metals.
Potassium is lower in the group.
Potassium has:
- more occupied shells
- larger atomic radius
- greater shielding
- lower first ionization energy
Its outer electron is easier to remove.
Therefore:
potassium is more reactive than lithium.
Worked Example 2
Which is more reactive: chlorine or iodine?
Both are Group 17 halogens.
Chlorine is higher in the group.
Chlorine has:
- fewer occupied shells
- smaller atomic radius
- less shielding
Its nucleus attracts an incoming electron more effectively.
Therefore:
chlorine is more reactive than iodine.
Worked Example 3
Will bromine displace iodide ions?
Group 17 reactivity:
F > Cl > Br > I
Bromine is more reactive than iodine.
Therefore:
yes
For example:
Br₂ + 2KI → 2KBr + I₂
Worked Example 4
Will iodine displace chloride ions?
Chlorine is more reactive than iodine.
Therefore, iodine cannot displace chloride ions under the usual halogen-displacement conditions.
Prediction:
no reaction
Worked Example 5
Element X is below sodium in Group 1.
Predict its ion and relative reactivity.
Group 1:
X⁺
Because X is lower than sodium:
- radius is larger
- shielding is greater
- outer electron is easier to remove
Therefore:
X should be more reactive than sodium.
Worked Example 6
Element Y is above bromine in Group 17.
Predict its ion and relative reactivity.
Group 17:
Y⁻
Because Y is higher than bromine:
- radius is smaller
- shielding is lower
- attraction for an incoming electron is stronger
Therefore:
Y should be more reactive than bromine.
Worked Example 7
Element X has electron arrangement:
2,8,2
Element Y has:
2,8,7
Predict their reaction.
X has two valence electrons and is likely to form:
X²⁺
Y has seven valence electrons and is likely to form:
Y⁻
Two Y⁻ ions are required to balance X²⁺.
Therefore, the expected ionic compound is:
XY₂
Worked Example 8
An unknown element is:
- metallic
- in Group 1
- below potassium
Predict its behaviour with water.
Because Group 1 reactivity increases down the group, the element should react very readily with water.
We would predict products of the general form:
metal hydroxide + hydrogen
and expect the reaction to be more vigorous than potassium's, assuming comparable conditions.
Common Mistakes
Mistake 1: Assuming reactivity always increases down a group
It depends on the group.
Group 1:
increases
Group 17:
decreases
Mistake 2: Explaining reactivity only by saying "it is lower in the group"
Position describes the trend but does not explain it.
Use:
shells → radius → shielding → attraction → electron gain/loss → reactivity
Mistake 3: Saying Group 1 gains electrons
Group 1 metals:
lose electrons
and form:
+1 ions
Mistake 4: Saying halogens lose electrons
Halogens:
gain electrons
and form:
−1 ions
Mistake 5: Assuming noble gases are highly reactive because they are on the right side
Noble gases have complete outer electron shells and are generally very unreactive.
Mistake 6: Treating all transition metals as having the same reactivity
Transition-metal reactivity varies considerably.
Mistake 7: Confusing ion charge with reactivity
Lithium and potassium both form +1 ions.
Their ion charges are the same, but their reactivities are different.
Mistake 8: Making a prediction without evidence
A strong chemistry answer should state both:
what is predicted
and:
why the periodic trend supports that prediction.
Did You Know?
The periodic table became one of chemistry's most important tools because it could be used to make successful predictions about elements that had not yet been discovered.
Dmitri Mendeleev left gaps in his periodic table where he believed undiscovered elements belonged.
He predicted several of their properties by examining trends among surrounding elements.
Later discoveries showed that many of these predictions were remarkably useful.
Modern atomic theory now explains these patterns in terms of electron arrangements and atomic structure.
Key Terms
Ion: A particle with an overall electrical charge.
Cation: A positively charged ion formed by electron loss.
Anion: A negatively charged ion formed by electron gain.
Valence electron: An electron in the outermost occupied shell.
Reactivity: The tendency of a substance to undergo chemical change.
Atomic radius: A measure of atomic size.
Shielding: Reduction in effective nuclear attraction caused by inner electrons.
Ionization energy: Energy required to remove an electron from an isolated gaseous atom.
Electronegativity: An atom's tendency to attract bonding electrons.
Displacement reaction: A reaction in which a more reactive element replaces a less reactive element.
Periodic trend: A repeating pattern in element properties across the periodic table.
Periodicity: The repeating pattern of physical and chemical properties associated with atomic structure.
Key Reactivity Patterns
Group 1
One valence electron.
Forms:
+1 ions
Reacts by:
losing an electron
Down the group:
reactivity increases
because electron removal becomes easier.
Group 2
Two valence electrons.
Forms:
+2 ions
Reacts mainly through:
electron loss
Group 17
Seven valence electrons.
Forms:
−1 ions
Reacts by:
gaining an electron
Down the group:
reactivity decreases
because attraction for an incoming electron becomes weaker.
Group 18
Complete outer shell.
Generally:
very unreactive
No common simple ion.
Transition Metals
Usually form:
positive ions
May have:
multiple oxidation states
Reactivity varies between different metals.
Key Takeaways
- The periodic table can be used to predict how unfamiliar elements are likely to react.
- Chemical reactivity is strongly connected to electron arrangement.
- Metals generally react by losing electrons and forming positive ions.
- Reactive non-metals often gain electrons and form negative ions.
- Group 1 metals form +1 ions and become more reactive down the group.
- Group 2 metals commonly form +2 ions.
- Group 17 halogens form −1 ions and become less reactive down the group.
- Group 18 noble gases have complete outer shells and are generally very unreactive.
- Transition metals form positive ions but often have variable oxidation states and do not follow one simple reactivity trend.
- Atomic radius generally increases down a group.
- Electron shielding generally increases down a group.
- First ionization energy helps predict how easily an atom can lose an electron.
- Electronegativity helps describe how strongly atoms attract bonding electrons.
- A more reactive halogen can displace a less reactive halogen from a halide compound.
- More reactive metals can displace less reactive metals in appropriate displacement reactions.
- Ion charge and reactivity are related to electron structure but are not the same property.
- Reaction outcomes can often be predicted by combining group membership, ion charges, and reactivity trends.
- Strong explanations should use evidence rather than simply state a memorized trend.
- For unfamiliar problems, use:
periodic position → electron arrangement → atomic radius and shielding → tendency to gain or lose electrons → ion formation → relative reactivity → predicted reaction outcome.