Periodic Table Families & Element Types

Metals

Groups 1–12, parts of 13–16, Periods 2–7)

Metals Characteristics

• Good conductors of heat & electricity

• Malleable (can be hammered) & ductile (can be drawn into wires)

• Usually solid at room temp (except Hg)

• Tend to lose electrons → form cations and have high melting and boiling points and are shiny

Examples of Metals

Na, Fe, Al, Cu, Au

Nonmetals Location

Upper right side of the periodic table (right of the staircase line)

Nonmetals Characteristics

• Poor conductors (insulators)

• Brittle if solid

• Many are gases at room temp

• Tend to gain electrons → form anions

Examples of nonmetals

Hydrogen (H), Carbon (C), Nitrogen (N), Oxygen (O), Phosphorus (P), Sulfur (S), Selenium (Se), 

The Halogens Fluorine Chlorine Bromine Iodine 

The Noble Gases Helium , Neon , Argon Krypton, Xenon,  Radon

Metalloids Location

Along the staircase line between metals and nonmetals

Metalloids Characteristics

• Have properties of both metals and nonmetals

• Semiconductors (useful in electronics)

• shiny, metallic luster but are brittle, unlike most metals. 

Metalloids Examples

1. Boron (B)

2. Silicon (Si)

3. Germanium (Ge)

4. Arsenic (As)

5. Antimony (Sb)

6. Tellurium (Te)

Alkali Metals Location

Group 1 not Hydrogen

Alkali Metals Characteristics

• Extremely reactive (especially with water)

• Soft, silvery metals so soft they can be cut with a knife. 

• 1 valence electron → form +1 ions

• Not found free in nature

• often stored in oil to prevent reactions

• have relatively low melting points compared to other metals. 

Examples of Alkali Metals

lithium (Li), sodium (Na),

potassium (K), rubidium (Rb),

cesium (Cs), and francium (Fr)

Alkaline Earth Metals location

Group 2

Alkaline Earth Metals Characteristics:

• Reactive, but less than alkali metals

• 2 valence electrons → form +2 ions

• Each atom has two valence electrons, which are easily removed to achieve a stable, noble gas configuration

• Harder than alkali metals

• They are shiny, silvery-white metals

• relatively low densities, melting points, and boiling points

Examples of Alkaline Earth Metals

beryllium, magnesium, calcium,

strontium, barium, and radium.

Transition Metals Location

(Groups 3–12)

Transition Metals Characteristics

• Typical metals (good conductors, malleable, ductile)

• Oxidation States:

  They can form ions with variable charges, unlike many other elements that have a single charge. 

• Colored Compounds:

  Their compounds are often brightly colored due to d-d electronic transitions

• high melting and boiling points

Inner Transition Metals

• Lanthanides: Period 6 (atomic #57–71) → shiny, reactive metals

• Actinides: Period 7 (atomic #89–103) → mostly radioactive
  Characteristics:

• Many are synthetic (especially actinides)

• Similar metallic properties

Halogens location

Group 17

Halogens Characteristics:

• Very reactive nonmetals

• 7 valence electrons → form -1 ions

• React with metals to form salts

• They are non-metals and are typically toxic. 

Diatomic Elements (Exist naturally as pairs)

Mnemonic: “Have No Fear Of Ice Cold Beer”

Diatomic Elements Characteristics

• Form molecules with two atoms (e.g., O₂, Cl₂)

• Mostly nonmetals

• Important for natural reactions (e.g., oxygen in respiration)

Examples of Diatomic Elements

H₂, N₂, F₂, O₂, I₂, Cl₂, Br₂

Noble Gases Location 

Group 18

Noble Gases Characteristics

• Very stable & unreactive (full valence shell)

• Colorless, odorless gases

• Used in lighting & balloons

When comparing elements in the same column, which factor is more important for determining attractive force: distance or number of protons?

Distance to the nucleus is the dominant factor because added energy levels increase distance and shielding, weakening attraction more than extra protons strengthen it.

What happens to the attractive force between the nucleus and outer electrons as you go down a group?

The attractive force decreases because electrons are farther from the nucleus and more shielded by inner electrons.

What is the mathematical relationship between distance (d) and attractive force (F)?

More protons = stronger positive charge = stronger attraction to electrons.

Why does increasing the number of protons not always lead to a stronger attraction in the outer shell?

Because increased distance and electron shielding can reduce the attractive force more than the extra protons increase it.