Introductory General Chemistry · Unit 05

Periodic Table & Trends

Start here if periodic trends still feel like arrows to memorize. This unit turns the table into a prediction tool — building from electron configuration and showing why these patterns matter for nomenclature and bonding.

What you'll learn

Read periods, groups, regions, and valence electron patterns. Predict atomic radius, ionization energy, and electronegativity trends. Compare atom size with ion size for cations and anions. Practice trend ranking with worked examples and interactive tools.

5.1 Start Here: How the Periodic Table Is Organized

The periodic table arranges all known elements in order of increasing atomic number (number of protons). It is set up so that elements with similar properties land in the same column.

Start here with the two directions that matter most:

Periodic Table Organization: Periods and Groups
Direction	Called	How many?	What changes as you move?
↔ Horizontal row	Period	7 periods	Elements in the same period have the same outermost energy level. Atomic number increases by 1 each step from left to right.
↕ Vertical column	Group (or Family)	18 groups	Elements in the same group have similar valence electron patterns. Moving down a group adds one energy level each row.

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Memory hook

"Period = row, Group = column." Period 2 elements all use the 2nd energy level as their outermost shell.
Group 1 elements all have 1 valence electron.
Do not mix up period and group. Getting those mixed up sends trend answers in the wrong direction.

5.2 Regions of the Periodic Table

Start by reading the table in three visible regions: left-side metals, staircase metalloids, and right-side nonmetals, plus a few important named families. Get these clean before you move to trend arrows — the regions tell you where each trend is strongest.

Periodic Table Regions and Element Types
Region	Location	Examples	Key traits
Alkali metals (Group 1)	Far left column	Li, Na, K	1 valence e⁻, very reactive, soft metals
Alkaline earth metals (Group 2)	Second column	Be, Mg, Ca	2 valence e⁻, reactive metals
Transition metals	Middle block (Groups 3–12)	Fe, Cu, Zn	Typical metals that often form more than one ion charge
Metalloids	Staircase border	Si, Ge, As	Properties between metals and nonmetals
Nonmetals	Upper right	C, N, O, S	Poor conductors, gain electrons
Halogens (Group 17)	Second to last column	F, Cl, Br	7 valence e⁻, very reactive nonmetals
Noble gases (Group 18)	Far right column	He, Ne, Ar	8 valence e⁻ (He: 2), almost no reactivity

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The staircase line separates metals (left) from nonmetals (right).
Elements touching the staircase are metalloids with mixed properties.
Silicon (Si) is the classic example: it conducts electricity, but only a little.

Here is how all three big trends map onto the table at once. Keep this as a reference as you work through 5.4–5.6.

One Map for the Three Big Trends

Atomic radius

Across a period: smaller to the right

Down a group: larger downward

Largest area: bottom-left

Ionization energy

Across a period: higher to the right

Down a group: lower downward

Highest area: top-right

Electronegativity

Across a period: higher to the right

Down a group: lower downward

Highest area: top-right, especially F

Quick anchor: radius points toward the bottom-left. Ionization energy and electronegativity point toward the top-right.

5.3 Valence Electrons and the Group Number

Valence electrons are the electrons in the outermost occupied energy level. They are the electrons most involved in chemical bonding. For main-group elements only (Groups 1, 2, and 13–18), the group number tells you the number of valence electrons. This is one of the most useful shortcuts in the whole course.

Shortcut for main-group elements Group 1 → 1 valence electron | Group 2 → 2 valence electrons
Group 13 → 3 Group 14 → 4 Group 15 → 5 Group 16 → 6
Group 17 → 7 Group 18 → 8 (He has 2)

Group Number and Valence Electron Examples
Element	Group	Valence e⁻
Na (sodium)	1	1
Mg (magnesium)	2	2
Al (aluminum)	13	3
C (carbon)	14	4
N (nitrogen)	15	5
O (oxygen)	16	6
Cl (chlorine)	17	7
Ar (argon)	18	8

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Elements in the same group have the same number of valence electrons, which is why they behave similarly.
Do not use this shortcut for transition metals in Groups 3–12.
Sodium (Na) and potassium (K) are both in Group 1, both have 1 valence electron, and both react violently with water.

5.4 Atomic Radius

The atomic radius is roughly half the distance between two bonded atoms of the same element. It tells us how big an atom is. Start here with the big idea: more shells usually means bigger, and more nuclear pull in the same shell usually means smaller.

Across a period (left → right): radius gets SMALLER

As you move right across a period, the number of protons increases, but the electrons are all in the same energy level. More protons pull the electrons inward more strongly. The atom shrinks.

Example — Period 3: Na (186 pm) → Mg (160 pm) → Al (143 pm) → Si (111 pm) → P (106 pm) → Cl (99 pm)

Down a group (top → bottom): radius gets LARGER

As you move down a group, each new period adds a whole new energy level. More energy levels mean electrons are farther from the nucleus. The atom grows.

Example — Group 1: Li (152 pm) → Na (186 pm) → K (227 pm) → Rb (248 pm) → Cs (265 pm)

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The biggest atoms are in the bottom-left corner (Cs, Fr).
The smallest atoms are in the top-right corner (He, F).
Picture the periodic table as a hillside: atoms get smaller as you climb toward the upper right.

Ready to test yourself on atomic radius? Practice Unit 05 Problems →

5.5 Ionization Energy

Ionization energy (IE) is the energy needed to remove one electron from a neutral atom in the gas phase. A high IE means the atom holds that electron tightly. A low IE means the electron is easier to remove. If you get stuck, ask one question: how hard is this atom fighting to keep its electron?

Across a period (left → right): ionization energy usually increases

As the number of protons increases, the nucleus pulls electrons more strongly, so it takes more energy to remove one. Noble gases are usually the highest in their period.

Down a group (top → bottom): ionization energy usually decreases

The outer electrons are farther from the nucleus and more shielded by inner electrons, so they are easier to remove.

Example — Group 1: Li (520 kJ/mol) → Na (496) → K (419) → Cs (376)

A few small exceptions occur near Groups 13 and 16 in the IE trend. Work through the examples below before worrying about those.

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IE trends are exactly opposite to atomic radius trends.
Smaller atom = harder to remove an electron.
Larger atom = easier to remove an electron.
Common mistake: ranking IE the same way as radius. They go opposite directions.

5.6 Electronegativity

Electronegativity (EN) measures how strongly an atom pulls shared electrons toward itself in a bond. Fluorine (F) is the most electronegative element at 3.98.

Electronegativity is not the same as ionization energy. Ionization energy is about removing an electron; electronegativity is about pulling on shared electrons in a bond.

Across a period (left → right): EN increases

More protons → stronger pull on shared electrons.

Down a group (top → bottom): EN decreases

Larger atom → nucleus is farther from bonding electrons → weaker pull.

Using ΔEN as a classroom guideline for bond type

ΔEN = |EN of element A − EN of element B|

About 0.0-0.4 → usually nonpolar covalent

About 0.5-1.7 → usually polar covalent

Greater than about 1.7 → usually ionic

Electronegativity Difference and Bond Type Examples
Bond	EN values	ΔEN	Bond type
H–Cl	2.20 & 3.16	0.96	Polar covalent
H–H	2.20 & 2.20	0.00	Nonpolar covalent
Na–Cl	0.93 & 3.16	2.23	Ionic
C–O	2.55 & 3.44	0.89	Polar covalent
K–F	0.82 & 3.98	3.16	Ionic

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EN follows the same direction as ionization energy: both increase toward the upper right of the table (excluding noble gases).
F is highest, Fr and Cs are lowest.
Do not mix up "holds its own electron" with "pulls shared electrons." That is the IE vs EN difference.

Where This Shows Up Next

In Unit 10 Bonding, you will use ΔEN to classify bonds as nonpolar covalent, polar covalent, or ionic.

In Unit 10 IMFs, bond polarity helps determine molecular polarity, which then helps predict intermolecular forces.

5.7 Ion Formation and Ionic Radius

When an atom gains or loses electrons to form an ion, its size changes significantly. Here is the size rule: losing electrons usually shrinks the particle, while gaining electrons usually makes it larger.

Metals LOSE electrons → cations (positive ions) → SMALLER than the atom

Na atom: 186 pm → Na⁺ ion: 102 pm (45% smaller!)

Reason: Fewer electrons, same number of protons. The remaining electrons are pulled in tighter.

Nonmetals GAIN electrons → anions (negative ions) → LARGER than the atom

Cl atom: 99 pm → Cl⁻ ion: 181 pm (83% bigger!)

Reason: More electrons, same number of protons. Electrons repel each other and spread out.

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Isoelectronic series

Ions with the same electron count can be compared by proton count.
More protons = smaller ion.
N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺, Al³⁺ all have 10 electrons.
Al³⁺ (13 protons) is the smallest.
N³⁻ (7 protons) is the largest.

These four rows show what actually happens to size when an atom becomes an ion — use them to check your reasoning on the two steps above.

Concrete Ion Size Comparisons
Species	What changed?	Radius	What the size change tells you
Na	Neutral atom	186 pm	Still has its 3rd energy level.
Na⁺	Lost 1 electron	102 pm	Losing the 3s electron removes the outer level, so the ion gets much smaller.
Cl	Neutral atom	99 pm	Has 17 electrons balancing 17 protons.
Cl⁻	Gained 1 electron	181 pm	Adding an electron increases repulsion in the valence shell, so the ion gets larger.

Quick Size Rule Cation = lost electrons = usually smaller than the atom
Anion = gained electrons = usually larger than the atom

Use these like a guided class check. Identify the family or predict the trend before you reveal the reasoning. That is the fastest way to clean up the usual mistakes before nomenclature and bonding build on these patterns. For more reps, use the Unit 05 Practice page.

Periodic Table Family Classifier

Highlighted element

Choose one family or region before you check your answer.

How To Read It

Which family does this highlighted element belong to?

Start with the highlighted element's column. Families are vertical groups on the periodic table.

Scenario: Use the highlighted element's position and choose the best family or region before you check.

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Quick reminder

Group 1 = alkali metals.
Group 2 = alkaline earth metals.
Group 17 = halogens.
Group 18 = noble gases.

Choose one answer, then click Check.

Alkali metal Alkaline earth Transition metal Post-transition metal Metalloid Nonmetal Halogen Noble gas Lanthanide Actinide

Periodic Trend Predictor

Read This First

Atomic Radius

Which atom has the larger atomic radius?

Decide whether this is an across-a-period comparison or a down-a-group comparison before you choose.

Compare: Na vs Cl

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Quick reminder

Atomic radius increases down and to the left.
Ionization energy and electronegativity increase up and to the right.

Choose one element, then click Check.

Ion Size Chooser

Read This First

Which species is larger?

Start by asking whether the ion lost electrons to form a cation or gained electrons to form an anion.

Compare: Na vs Na⁺

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Quick reminder

Cation = lost electrons = usually smaller.
Anion = gained electrons = usually larger.

Choose one species, then click Check.

Example 1 · Identifying Period, Group, and Region

Problem: Use sulfur's position to identify its place and behavior on the periodic table.

Given: Element sulfur (S), atomic number 16

Find: period, group, number of valence electrons, and region of the periodic table

Step 1 — Find sulfur on the periodic table

Sulfur (S) is in Period 3 and Group 16.

Step 2 — Use the group to find valence electrons

Sulfur is a main-group element in Group 16, so it has 6 valence electrons.

Step 3 — Identify the region

Sulfur is on the right side of the table and is not on the staircase line. It is a nonmetal.

Step 4 — Check the answer with the electron configuration

S: 1s² 2s² 2p⁶ 3s² 3p⁴. The outermost occupied energy level is 3, and it contains 6 electrons.

Sulfur: Period 3 · Group 16 · 6 valence electrons · Nonmetal

Example 2 · Ranking Atomic Radii

Problem: Rank the atoms from smallest to largest atomic radius.

Given: Na, Mg, K, Cl

Find: smallest → largest radius

Step 1 — Locate each element on the table

Example 2 Element Positions for Atomic Radius Ranking
Element	Period	Group
Na	3	1
Mg	3	2
K	4	1
Cl	3	17

Step 2 — Apply the period trend

Na, Mg, and Cl are all in Period 3. Across a period, atomic radius gets smaller from left to right, so the Period 3 order is Na > Mg > Cl.

Step 3 — Apply the group trend

K is below Na in Group 1. Atomic radius gets larger as you move down a group, so K is larger than Na.

Step 4 — State the final ranking

Cl < Mg < Na < K

Example 3 · Using Electronegativity to Classify a Bond

Problem: Use electronegativity values to classify the bond type.

Given: Bond between potassium (K) and fluorine (F)

Find: bond type

Step 1 — Look up the EN values

K: 0.82 | F: 3.98 (F is the most electronegative element)

Step 2 — Calculate ΔEN

ΔEN = |EN_K − EN_F| = |0.82 − 3.98| = 3.16

Step 3 — Apply the bond-type guideline

3.16 > 1.7, so this falls in the ionic range of the classroom guideline.

This makes sense: K is a metal in the far bottom-left, F is a nonmetal in the top-right. Large EN differences always occur between metals and nonmetals.

In introductory chemistry, these ΔEN cutoffs are used as a guideline for classifying bond type.

Step 4 — State the final answer

KF is classified as ionic because the electronegativity difference is very large.

Example 4 · Comparing Atom and Ion Size

Problem: Decide whether the atom or ion is larger.

Given: Sodium atom (Na) and sodium ion (Na⁺)

Find: which species is larger

Step 1 — Write out what changed

Na is a Group 1 metal. It loses 1 electron to form Na⁺.

Na atom: 11 protons, 11 electrons → Na⁺ ion: 11 protons, 10 electrons

Step 2 — Track what losing an electron does to size

Na loses its entire outermost energy level (3s¹). The remaining 10 electrons are all in level 2, much closer to the nucleus.

Na atom: 186 pm → Na⁺ ion: 102 pm

Step 3 — State the conclusion

Na atom is larger than Na⁺ because losing the outermost electron makes the ion smaller.

Example 5 · Finding Valence Electrons from Position

Problem: Find valence electrons without writing a full electron configuration first.

Given: Phosphorus (P), atomic number 15

Find: number of valence electrons

Step 1 — Find P on the periodic table

P is in Period 3, Group 15.

Step 2 — Apply the main-group shortcut

For main-group elements, the last digit of the group number gives the number of valence electrons.

Group 15 → 5 valence electrons

Step 3 — Verify with the electron configuration

P: 1s² 2s² 2p⁶ 3s² 3p³ → 3s² + 3p³ = 5 electrons in level 3 ✓

Introductory General Chemistry · Unit 05 · Periodic Table & Trends