Periodic Trends

📖 Learn

Periodic trends are predictable patterns in the properties of elements that occur across periods (rows) and down groups (columns) of the periodic table. These trends arise from the arrangement of electrons in atoms and help us predict how elements will behave chemically.

Why Do Periodic Trends Exist?

Periodic trends are governed by three main factors:

Nuclear Charge: The number of protons in the nucleus (atomic number) - more protons means stronger pull on electrons
Electron Shielding: Inner electrons shield outer electrons from the full attraction of the nucleus
Number of Energy Levels: More energy levels mean electrons are farther from the nucleus

Effective Nuclear Charge (Zeff)

Effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom. It equals the total nuclear charge minus the shielding effect of inner electrons: Zeff = Z - S (where Z is atomic number and S is shielding).

Major Periodic Trends

1. Atomic Radius

Atomic Radius

The atomic radius is half the distance between the nuclei of two bonded atoms of the same element.

Direction	Trend	Reason
Across a period (left to right)	Decreases	More protons pull electrons closer; same energy level, increasing Zeff
Down a group	Increases	Additional energy levels make atoms larger despite more protons

2. Ionization Energy

Ionization Energy

Ionization energy (IE) is the energy required to remove an electron from a gaseous atom. First ionization energy removes the first electron; second ionization energy removes the second, and so on.

Direction	Trend	Reason
Across a period (left to right)	Generally increases	Higher Zeff holds electrons more tightly; smaller radius
Down a group	Decreases	Valence electrons farther from nucleus; easier to remove

Ionization Energy Exceptions

There are small dips in ionization energy between Groups 2 and 13 (Be to B, Mg to Al) because p electrons are slightly easier to remove than s electrons. Similarly, there's a dip between Groups 15 and 16 (N to O, P to S) because paired electrons repel each other and are easier to remove.

3. Electronegativity

Electronegativity

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. The Pauling scale is most commonly used, where fluorine has the highest value (4.0).

Direction	Trend	Reason
Across a period (left to right)	Increases	Higher Zeff attracts bonding electrons more strongly
Down a group	Decreases	Valence electrons farther from nucleus; weaker attraction

4. Electron Affinity

Electron Affinity

Electron affinity is the energy change when an electron is added to a gaseous atom. A more negative value indicates greater energy release and a more favorable process.

Direction	Trend	Reason
Across a period (left to right)	Generally becomes more negative	Higher Zeff makes adding electron more favorable
Down a group	Generally becomes less negative	Larger atoms have weaker attraction for added electrons

5. Metallic Character

Direction	Trend	Reason
Across a period (left to right)	Decreases	Atoms less willing to lose electrons; nonmetallic behavior increases
Down a group	Increases	Valence electrons more easily lost; metallic properties stronger

Summary of Trends

Property	Across Period (L to R)	Down Group
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electronegativity	Increases	Decreases
Electron Affinity	More negative	Less negative
Metallic Character	Decreases	Increases

Memory Tip

Notice that atomic radius has the opposite trend from ionization energy and electronegativity. This makes sense: smaller atoms hold their electrons more tightly (higher IE) and attract bonding electrons more strongly (higher electronegativity).

SAT/ACT Connection

Science sections often present data tables or graphs showing element properties and ask you to identify trends or predict values for unlisted elements. Understanding these trends allows you to make predictions even for elements you haven't studied directly.

💡 Examples

Example 1: Comparing Atomic Radii

Problem: Arrange the following elements in order of increasing atomic radius: Na, Mg, K, Ca.

Solution:

Step 1: Identify positions in the periodic table:

Na: Period 3, Group 1
Mg: Period 3, Group 2
K: Period 4, Group 1
Ca: Period 4, Group 2

Step 2: Apply trends: Atomic radius decreases across a period and increases down a group.

Step 3: Compare within each period: Mg < Na (Period 3); Ca < K (Period 4)

Step 4: Compare between periods: Period 3 elements are smaller than Period 4 elements.

Answer: Mg < Na < Ca < K (increasing atomic radius)

Example 2: Predicting Ionization Energy

Problem: Which element has the highest first ionization energy: Li, Na, or Ne?

Solution:

Step 1: Locate elements: Li (Period 2, Group 1), Na (Period 3, Group 1), Ne (Period 2, Group 18)

Step 2: Compare Li and Na: Both in Group 1, Na is below Li, so Na has lower IE.

Step 3: Compare Li and Ne: Both in Period 2, Ne is further right, so Ne has higher IE.

Step 4: Noble gases have very high ionization energies due to stable electron configurations.

Answer: Ne has the highest first ionization energy.

Example 3: Electronegativity Differences

Problem: Between oxygen (O) and sulfur (S), which is more electronegative? Explain using periodic trends.

Solution:

Step 1: Both O and S are in Group 16 (chalcogens).

Step 2: O is in Period 2, S is in Period 3.

Step 3: Electronegativity decreases down a group.

Step 4: Therefore, O (higher in the group) is more electronegative than S.

Answer: Oxygen is more electronegative than sulfur because electronegativity decreases as you go down a group. Oxygen's valence electrons are closer to the nucleus, allowing stronger attraction to bonding electrons.

Example 4: Ion Size Comparison

Problem: Which is larger: Na+ or Na? Explain.

Solution:

Step 1: Na (sodium atom) has 11 protons and 11 electrons.

Step 2: Na+ has 11 protons but only 10 electrons (lost one electron).

Step 3: When Na loses an electron to become Na+:

The entire outer energy level (3s) is removed
Electron-electron repulsion decreases
The 11 protons now pull on only 10 electrons, increasing Zeff

Answer: Na (neutral atom) is larger than Na+. Cations are always smaller than their parent atoms because they have fewer electrons and often lose an entire energy level.

Example 5: Using Multiple Trends

Problem: Predict which element would be the best reducing agent: Na, Mg, or Al.

Solution:

Step 1: A reducing agent loses electrons easily. This correlates with:

Low ionization energy
Low electronegativity
High metallic character

Step 2: All three are in Period 3. Moving left to right: Na, Mg, Al

Step 3: Ionization energy increases left to right, so Na has the lowest IE.

Step 4: Metallic character decreases left to right, so Na is the most metallic.

Answer: Na (sodium) is the best reducing agent because it has the lowest ionization energy and loses its valence electron most easily.

✏️ Practice

1. Which element has the largest atomic radius?

A) Li

B) Na

C) K

D) Rb

2. As you move from left to right across Period 3, what happens to electronegativity?

A) It increases

B) It decreases

C) It stays the same

D) It first increases, then decreases

3. Which of the following has the highest first ionization energy?

A) Sodium (Na)

B) Magnesium (Mg)

C) Aluminum (Al)

D) Chlorine (Cl)

4. Which statement correctly describes the trend for metallic character?

A) Increases across a period and down a group

B) Decreases across a period and down a group

C) Decreases across a period and increases down a group

D) Increases across a period and decreases down a group

5. Which ion is larger: Cl- or K+?

A) Cl- because anions are larger than cations with the same number of electrons

B) K+ because it has more protons

C) They are the same size because they are isoelectronic

D) K+ because potassium is a larger atom than chlorine

6. What causes ionization energy to generally increase across a period?

A) Increasing atomic radius

B) Increasing effective nuclear charge

C) Decreasing number of protons

D) Increasing electron shielding

7. Arrange the following in order of increasing electronegativity: C, N, O, F

A) F < O < N < C

B) C < N < O < F

C) N < C < F < O

D) O < F < C < N

8. Which element would have the most negative electron affinity?

A) Lithium (Li)

B) Carbon (C)

C) Fluorine (F)

D) Neon (Ne)

9. Why does atomic radius increase down a group?

A) The effective nuclear charge increases

B) Each element has an additional energy level of electrons

C) The number of protons decreases

D) Electron shielding decreases

10. Based on periodic trends, which element is most likely to form a +1 cation?

A) Oxygen

B) Neon

C) Cesium

D) Fluorine

Click to reveal answers

D - Rb (rubidium) is furthest down Group 1, giving it the largest atomic radius.
A - Electronegativity increases across a period due to increasing effective nuclear charge.
D - Chlorine has the highest ionization energy being furthest right in the period (excluding noble gases in the choices).
C - Metallic character decreases across (atoms hold electrons tighter) and increases down (electrons easier to lose).
A - Both have 18 electrons (isoelectronic), but Cl- has only 17 protons pulling on them while K+ has 19, making Cl- larger.
B - Increasing effective nuclear charge (Zeff) pulls electrons closer, making them harder to remove.
B - Electronegativity increases left to right across Period 2: C < N < O < F.
C - Fluorine has high electron affinity because adding one electron completes its octet. Noble gases (Ne) have positive electron affinity since they don't want additional electrons.
B - Additional energy levels place valence electrons further from the nucleus, increasing atomic size.
C - Cesium (Group 1, Period 6) has very low ionization energy and readily loses its single valence electron to form Cs+.

✅ Check Your Understanding

Question 1: Explain why noble gases have very high ionization energies but essentially zero electronegativity values.

Reveal Answer

Noble gases have complete outer electron shells (stable octets), making their electrons very difficult to remove - hence high ionization energy. However, electronegativity measures an atom's ability to attract shared electrons in a bond. Since noble gases rarely form bonds due to their stable configuration, they have no meaningful electronegativity value. They don't need or want additional electrons, so assigning them an electronegativity would be meaningless.

Question 2: Sulfur and chlorine are adjacent elements in Period 3. Predict three properties where chlorine would have a higher value than sulfur, and explain why.

Reveal Answer

Chlorine has higher values than sulfur for: (1) Ionization energy - Cl has one more proton creating stronger Zeff, holding electrons more tightly. (2) Electronegativity - Cl's higher Zeff makes it attract bonding electrons more strongly. (3) Electron affinity (more negative) - Cl only needs one electron to complete its octet, while S needs two, making Cl's electron addition more energetically favorable. All these trends result from chlorine being further right in the period with higher effective nuclear charge.

Question 3: A student claims that since potassium (K) is a larger atom than sodium (Na), potassium must have weaker metallic properties. Evaluate this claim.

Reveal Answer

The student's claim is incorrect. While K is indeed larger than Na (atomic radius increases down Group 1), metallic character actually increases down a group, not decreases. Potassium's larger size means its valence electron is farther from the nucleus and experiences more shielding. This makes K's valence electron easier to lose, which is a defining metallic property. Potassium is MORE metallic than sodium - it's more reactive with water, has lower ionization energy, and more readily forms K+ ions.

Question 4: Oxygen (O2-), fluorine (F-), sodium (Na+), and magnesium (Mg2+) are all isoelectronic (have 10 electrons each). Rank them in order of increasing size and explain your reasoning.

Reveal Answer

Order of increasing size: Mg2+ < Na+ < F- < O2-. All four species have exactly 10 electrons, so the size depends entirely on how many protons are pulling on those electrons. Mg2+ has 12 protons, Na+ has 11 protons, F- has 9 protons, and O2- has 8 protons. More protons mean a stronger pull on the same number of electrons, resulting in a smaller ion. Therefore, Mg2+ (most protons, smallest) < Na+ < F- < O2- (fewest protons, largest). This demonstrates that for isoelectronic species, size is inversely related to nuclear charge.

🚀 Next Steps

Review any concepts that felt challenging
Move on to the next lesson when ready
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