Periodic Table Study Notes

The Periodic Table

13.1 Arrangement of Elements

• The Periodic Table arranges elements by increasing proton number.
• Elements are organized into periods (horizontal rows) and groups (vertical columns).

Groups

• Groups are vertical columns, numbered I to VIII.
• They extend from top to bottom.

Periods

• Periods are horizontal rows, numbered 1 to 7.
• They run from left to right.

13.2 Periodic Trends

Metallic and Non-Metallic Properties

• Elements are classified as metals, non-metals, or metalloids.
• A diagonal line separates metals from non-metals.
• Metalloids have properties of both metals and non-metals.
• Silicon is a common metalloid used in computer chips.

Trends Across a Period

• Metals are on the left, non-metals on the right.
• Metallic properties decrease, and non-metallic properties increase from left to right.
• Oxides change from basic to amphoteric to acidic.
  • Example: Period 3 elements (Na, Mg, Al, Si, P, S, Cl, Ar).

Trends Down a Group

• Atomic size increases down a group.
• Outer shell electrons are further from the nucleus.
• Elements lose outer shell electrons more easily.
  • Example: Group I (Lithium, Sodium, Potassium).

Electronic Configuration

• Electronic configuration determines period and group number.

• Period number equals the number of electron shells.

  • Example: Elements in Period 2 have two electron shells, Period 3 have three.

• Group number equals the number of outer shell electrons.

  • Example: Group I elements have one outer shell electron, Group IV have four.

• Elements in the same group have similar chemical properties.

Group Number and Ion Charge

• Groups I, II, and III form positive ions (metals).
  • The charge of the ion is the same as the group number.
• Groups IV and V share electrons to form covalent bonds.
  • Elements have a maximum oxidation number equal to the group number.
• Groups VI and VII form negative ions (non-metals).
• Group VIII elements have stable electronic configurations and do not form compounds.

Trends in Groups

• Mendeleev arranged elements by properties, predicting new elements.
• Properties of elements in a group show trends.

13.3 Group I Elements: Alkali Metals

• Alkali metals are in Group I.
• They have similar properties.

Physical Properties

• Soft and easily cut.
• Low melting points.
• Low densities (Lithium, Sodium, and Potassium float on water).

Trends Down Group I

• Melting points decrease.
• Densities generally increase.

Chemical Properties

• Highly reactive, stored in oil.
• Each has one outer shell electron, easily lost to achieve noble gas configuration.
• Reactivity increases down the group due to increased atomic size and easier electron loss.

Reactions with Water

• Alkali metal + cold water → alkali + hydrogen
• Lithium: $2Li(s) + 2H<em>2O(l) \rightarrow 2LiOH(aq) + H</em>2(g)$
• Sodium: $2Na(s) + 2H<em>2O(l) \rightarrow 2NaOH(aq) + H</em>2(g)$
• Potassium: $2K(s) + 2H<em>2O(l) \rightarrow 2KOH(aq) + H</em>2(g)$

Reducing Agents

• Alkali metals are powerful reducing agents (lose electrons readily).
  • Example: $Li \rightarrow Li^+ + e^-$
• Reducing power increases down the group.

Ionic Compounds

• Form ionic compounds with similar chemical formulas.
• Soluble in water.
• Examples: Carbonates, Nitrates, Sulfates, Chlorides.

13.4 Group VII Elements: Halogens

• Halogens are in Group VII: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At).

Physical Properties

• Non-metals, diatomic covalent molecules (e.g., $F<em>2$, $Cl</em>2$, $Br<em>2$, $I</em>2$).
• Low melting and boiling points.
• Coloured.

Trends Down the Group

• Melting and boiling points increase.
• Density increases.
• Colours become darker.

Chemical Properties

• Reactive non-metals with seven outer shell electrons (need one more for a stable configuration).
• React with metals to form salts (halides).

Reactivity

• Reactivity decreases down the group.
• Harder for the nucleus to attract an electron as atomic size increases.
• Order of reactivity: Chlorine > Bromine > Iodine

Displacement Reactions

• More reactive halogen displaces a less reactive one from its halide solution.
  • Example: $Cl<em>2(aq) + 2NaBr(aq) \rightarrow 2NaCl(aq) + Br</em>2(aq)$

Oxidizing Agents

• Halogens are powerful oxidizing agents (gain electrons readily).
• $X_2 + 2e^- \rightarrow 2X^-$
• Fluorine is the strongest oxidizing agent; oxidizing power decreases down the group.

Redox Reactions

• Displacement reactions are redox reactions.
• Example: $Cl<em>2(aq) + 2Br^-(aq) \rightarrow 2Cl^-(aq) + Br</em>2(aq)$
  • Chlorine is reduced, Bromide ion is oxidized.

13.5 Transition Elements

• Located between Groups II and III.
• Examples: Chromium (Cr), Manganese (Mn), Iron (Fe), Copper (Cu).

Properties

• High melting points and densities.
• Form coloured compounds (different colours at different oxidation numbers).
• Good catalysts.
• Variable oxidation numbers.

Melting Points and Densities

• Higher than Group I and II metals.

Coloured Compounds

• Examples: Chromium(III) chloride, Iron(II) sulfate, Manganese(IV) oxide, Copper(I) oxide

Catalysts

• Used in many reactions in the lab and industry.
• Examples: Iron (Haber process), Nickel (margarine manufacture)

Variable Oxidation Numbers

• Unlike Group I and II metals, transition metals form ions with different oxidation numbers.
• Examples: Chromium (+3, +6), Manganese (+2, +4, +7), Iron (+2, +3), Copper (+1, +2)

13.6 Noble Gases

• Group VIII elements: Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn).

Properties

• Non-metals.
• Monatomic.
• Colourless gases at room temperature.
• Low melting and boiling points.
• Insoluble in water.
• Unreactive.

Unreactive Nature

• Full outer shell electron configurations (Helium has 2, others have 8).
• Do not lose, gain, or share electrons.
• Used to provide inert atmospheres (e.g., Argon in tungsten bulbs).