Unit 2 Review- Patterns of the Periodic Table
What This Unit Actually Covers
Unit 2 on the periodic table isn't about memorizing every element. It's about understanding why elements behave the way they do. Once you get the patterns, you can predict chemical behavior without looking anything up. That's the whole point.
If you're struggling, it's probably because you're trying to memorize instead of recognizing patterns. Stop that. Start seeing the structure.
The Table's Basic Structure
The periodic table arranges elements by atomic number (protons) in horizontal rows called periods. There are 7 periods. Each period represents a new electron shell being filled.
Vertical columns are groups (or families). Elements in the same group have the same number of valence electrons. This shared electron count is why they react similarly.
Quick Reference
- Period = horizontal row (left to right)
- Group = vertical column (top to bottom)
- Atomic number = number of protons
- Valence electrons = electrons in the outermost shell
The Three Big Categories
Everything on the periodic table falls into one of these:
Metals
Left side. They're shiny, conductive, malleable, and they lose electrons during reactions. Most elements are metals.
Nonmetals
Upper right corner (excluding noble gases). They're brittle, poor conductors, and they gain or share electrons during reactions.
Metalloids
The staircase line between metals and nonmetals. Boron, silicon, germanium, arsenic, antimony, tellurium, polonium. These have intermediate properties—they conduct electricity but poorly. Silicon is the big one for electronics.
Major Periodic Trends You Need to Know
Trends run in two directions. Learn them. Know them. Questions will ask you to explain them.
Across a Period (Left to Right)
Atomic radius decreases. Shielding stays roughly the same, but nuclear charge increases, pulling electrons closer.
Ionization energy increases. Harder to remove an electron when the nucleus has a stronger pull.
Electronegativity increases. Elements crave electrons more as you move right.
Down a Group (Top to Bottom)
Atomic radius increases. You're adding electron shells, so the outer electrons are farther from the nucleus.
Ionization energy decreases. Outer electrons are farther away and more shielded, so they're easier to remove.
Electronegativity decreases. Same reason—outer electrons are farther and more shielded.
Summary Table
| Trend | Across Period (L → R) | Down Group (T → B) |
|---|---|---|
| Atomic Radius | Decreases | Increases |
| Ionization Energy | Increases | Decreases |
| Electronegativity | Increases | Decreases |
| Metallic Character | Decreases | Increases |
Why These Trends Exist
Two factors drive everything:
- Nuclear charge — more protons pull electrons closer
- Shielding — inner electrons block the nuclear pull on outer electrons
Across a period: nuclear charge increases, shielding stays constant. Net pull increases. Electrons get pulled in.
Down a group: nuclear charge increases, but you're adding entire shells of shielding. Shielding wins. Outer electrons feel less pull.
That's it. Two factors. Everything else is just application.
Ionization Energy Quirks
Ionization energy doesn't increase perfectly smoothly across a period. There are drops where you'd expect rises.
- Group 2 elements have higher ionization energy than Group 13 because removing an electron from a filled s-subshell is hard
- Group 15 elements have higher ionization energy than Group 16 because the half-filled p-subshell is stable
These exceptions show up on exams. Memorize them or understand why they happen. Understanding is faster.
Electronegativity and Bonding
Electronegativity measures how strongly an atom pulls on shared electrons in a bond.
Fluorine is the most electronegative element at 3.98. Cesium is the least at 0.79.
The difference between bonded atoms determines bond type:
- Difference < 0.4 = nonpolar covalent
- Difference 0.4–1.7 = polar covalent
- Difference > 1.7 = ionic
Getting Started: How to Answer Trend Questions
When an exam question asks you to explain a trend, follow this pattern:
- Identify the two elements or positions
- State the direction of the trend
- Cite the cause (increased nuclear charge, increased shielding, or distance from nucleus)
Example question: Why does oxygen have a higher ionization energy than sulfur?
Answer: Both are in Group 16, but sulfur is below oxygen in the group. Sulfur has more electron shells, so its valence electrons are farther from the nucleus and more shielded by inner electrons. The nuclear pull on the outer electrons is weaker, making sulfur easier to ionize. Therefore, oxygen has higher ionization energy.
Three sentences. Direct cause. No fluff.
Quick Memorization Tips
Don't write these out a hundred times. Use them once or twice and the patterns stick on their own.
- Atomic radius increases down and left
- Ionization energy and electronegativity increase up and right
- Metallic character increases down and left
- Remember: metals are on the left, big and soft—they lose electrons easily
What to Actually Study
Don't waste time re-reading the textbook. Practice these instead:
- Arrange elements by atomic radius given their positions
- Predict which element has higher ionization energy and explain why
- Identify bond types using electronegativity differences
- Label metals, nonmetals, and metalloids on a blank periodic table outline
If you can do those four things without hesitation, you're ready. If not, keep practicing until you can.