Representative Metals, Metalloids, and Nonmetals Study Guide

Classification and Periodicity of Representative Elements

• Representative Elements: Elements where the $s$ and $p$ orbitals are filling. These include groups 1, 2, and 12–18.

• Transition Elements: Elements where the $d$ orbitals (groups 3–11) are filling.

• Inner Transition Metals: Elements where the $f$ orbitals are filling.

• Group 12 Classification: Although they are technically in the d-block, the $d$ orbitals are filled by group 11. In group 12, the last electron enters an $s$ orbital, thus they qualify as representative metals.

• Metallic Character: Defined by an element's ability to lose its outer valence electrons. Characteristics include high thermal and electrical conductivity.

• Distribution of Representative Metals: There are 20 nonradioactive representative metals located in groups 1, 2, 3, 12, 13, 14, and 15.

• Metalloids: Elements with properties intermediate between metals and nonmetals, typically semiconductors. These include Boron ($B$), Silicon ($Si$), Germanium ($Ge$), Arsenic ($As$), Antimony ($Sb$), and Tellurium ($Te$).

• Nonmetals: Elements that tend to form anions or molecular compounds. They contrast with metals, which typically form cations and ionic compounds.

• Salts: Ionic compounds consisting of cations and anions, generally produced by the combination of a metal and a nonmetal.

• Passivation: The process by which some metals (e.g., $Be, Mg, Zn, Cd, Hg, Al, Sn, Pb$) react with air to form a nonreactive thin film of oxide or other compound, protecting the bulk metal from further reaction.

Group 1: The Alkali Metals

• Elements: Lithium ($Li$), Sodium ($Na$), Potassium ($K$), Rubidium ($Rb$), Cesium ($Cs$), and Francium ($Fr$).

• Etymology: The name "alkali metal" refers to the fact that these metals and their oxides react with water to form basic (alkaline) solutions.

• Atomic Properties:

  • Largest atomic radii in their periods.

  • Lowest first ionization energy in their periods.

  • Readily form stable cations with a charge of $1+$ by losing the single valence electron.

• Reactivity Trends: Reactivity increases with increasing atomic number due to the decrease in ionization energy.

• Physical Properties: Very soft solids. Lithium has the lowest density of any metal ($0.5123 \text{ g/cm}^3$, though usually rounded to $0.5 \text{ g/cm}^3$).

• Reaction with Water: React vigorously to form hydrogen gas and a basic solution of the metal hydroxide.

  • Example: $2Li(s) + 2H_2O(l) \rightarrow 2LiOH(aq) + H_2(g)$

• Storage: Due to high reactivity with moisture and oxygen, they are stored in sealed containers under mineral oil or kerosene.

• Applications:

  • Sodium vapor lamp streetlights use yellow light which penetrates fog better than white light.

  • Sodium is used to reduce compounds to produce metals like titanium ($Ti$) and zirconium ($Zr$).

  • Flame colors: Lithium (crimson), Sodium (bright yellow), Others (pale violet).

Group 2: The Alkaline Earth Metals

• Elements: Beryllium ($Be$), Magnesium ($Mg$), Calcium ($Ca$), Strontium ($Sr$), Barium ($Ba$), and Radium ($Ra$).

• Atomic Properties: Smaller atoms and higher first ionization energies than alkali metals in the same period due to increased nuclear charge.

• Reactivity: Less reactive than alkali metals but still highly reactive. Reactivity increases down the group. They form compounds in the $2+$ oxidation state.

• Bonding Character: Beryllium and magnesium form some compounds with covalent character due to higher ionization energies.

• Flame Colors: Calcium (red), Strontium (shades of red), Barium (green).

• Magnesium Specifics:

  • Malleable and ductile at high temperatures.

  • Passivated by a layer of magnesium oxycarbonate ($Mg_2(OH)_2CO_3$).

  • Highly flammable; burns with brilliant white light (used in flares).

  • Reaction with $CO_2$: $2Mg(s) + CO_2(g) \rightarrow 2MgO(s) + C(s)$. Note: $CO_2$ extinguishers cannot be used on magnesium fires.

• Solubility: Many common salts of alkaline earth metals are insoluble in water due to high lattice energies of divalent metal ion compounds.

Group 12: The Post-Transition Representative Metals

• Elements: Zinc ($Zn$), Cadmium ($Cd$), Mercury ($Hg$), and Copernicium ($Cn$).

• Electron Configuration: Each has two electrons in the outer shell ($ns^2$). Forming $2+$ cations results in a pseudo-noble gas electron configuration.

• Mercury ($Hg$): The only metal that is liquid at $25 \text{ } ^\text{o} C$. It exhibits a $1+$ oxidation state in the diatomic ion $Hg_2^{2+}$. It forms amalgams (alloys of mercury with other metals).

• Zinc ($Zn$): Most reactive in the group. Passivated by a basic carbonate coating ($Zn_2(OH)_2CO_3$). Used in dry cell batteries as an anode and for galvanizing steel (sacrificial anode).

• Mercury Reactivity: Nonreactive compared to hydrogen; does not displace $H_2$ from acids. Reacts with strong oxidizing acids like nitric acid:

  • $3Hg(l) + 8HNO_3(aq) \text{ (dilute)} \rightarrow 3Hg(NO_3)_2(aq) + 2NO(g) + 4H_2O(l)$

Group 13: Boron Family

• Elements: Boron ($B$), Aluminum ($Al$), Gallium ($Ga$), Indium ($In$), and Thallium ($Tl$).

• Bonding: Boron is a semiconducting metalloid with covalent binary compounds. The others are metals.

• Amphoterism: Aluminum and Gallium oxides and hydroxides react with both acids and bases.

  • Aluminum reaction with acid: $2Al(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2(g)$

  • Aluminum reaction with base: $2Al(s) + 2NaOH(aq) + 6H_2O(l) \rightarrow 2Na[Al(OH)_4](aq) + 3H_2(g)$

• Inert Pair Effect: Tendency of the two electrons in the outermost s orbital to remain unshared in compounds, leading to an oxidation state two lower than the group valence.

  • Visible in Thallium ($Tl^+$ is more stable than $Tl^{3+}$).

• Aluminum Properties: Lightweight, high conductivity, and corrosion-resistant due to passivation. Forms amalgams with mercury, which disrupts the protective oxide layer.

Group 14 and 15 Representative Metals

• Group 14 Metals: Tin ($Sn$), Lead ($Pb$), and Flerovium ($Fl$).

• Oxidation States: Stable divalent cations ($Sn^{2+}, Pb^{2+}$) due to the inert pair effect. They also form covalent compounds in the $4+$ state (e.g., $SnCl_4$ is a covalent liquid).

• Tin Allotropes:

  • White tin: Stable above $13.2 \text{ } ^\text{o} C$, malleable.

  • Gray tin: Stable at low temperatures, brittle, breaks into powder ("tin pest" or "tin disease").

• Lead Properties: Important in lead storage batteries; resistant to many acids but reacts with hot concentrated acid.

• Group 15 Metal: Bismuth ($Bi$). Least reactive representative metal. Exhibits $3+$ and rarely $5+$ (requires strong oxidizers) oxidation states.

Occurrence and Preparation of Representative Metals

• Natural Sources: Found in minerals: silicates, phosphates, carbonates, sulfates, or sulfides.

  • Aluminum: abundant in bauxite ($AlO(OH)$).

  • Magnesium: found in seawater (0.5 \text{% } MgCl_2).

• Electrolysis Preparation: Required for groups 1, 2, and aluminum.

  • Downs Cell (Sodium): Electrolysis of molten $NaCl$. Calcium chloride is added to lower the melting point from $801 \text{ } ^\text{o} C$ to $600 \text{ } ^\text{o} C$.

  • Hall-Héroult Cell (Aluminum): Purification of bauxite to $Al_2O_3$, dissolved in molten cryolite ($Na_3AlF_6$) and $CaF_2$. Electrolysis yields aluminum at the cathode.

• Chemical Reduction:

  • Pidgeon Process (Magnesium): $Si(s) + 2MgO(s) \rightarrow SiO_2(s) + 2Mg(g)$ at high temperature; forward reaction driven by removing magnesium vapor.

  • Zinc Production: Roasting sulfide to oxide, followed by reduction with carbon (coal).

  • Tin Production: Heating $SnO_2$ with carbon above $1000 \text{ } ^\text{o} C$.

Structure and General Properties of Metalloids

• Definition: Elements ($B, Si, Ge, As, Sb, Te$) that look metallic but are semiconductors because electrons are more tightly bound to nuclei than in metals.

• Silicon and Germanium: Crystallize in a diamond structure. Each atom is tetrahedral and $sp^3$ hybridized.

• Arsenic and Antimony: Layered structure with puckered sheets; atoms bonded to three neighbors within the sheet (similar to graphite).

• Boron Structure: Pure crystalline boron consists of $B_{12}$ icosahedra (20 faces, 12 corners). Bonds are approximately $176 \text{ pm}$ long.

• Silicon Preparation: Produced by high-temperature reduction of $SiO_2$ with carbon. Purified via Zone Refining, where impurities dissolve in a moving molten zone and are pushed to the end of a rod.

• Reactivity:

  • Silicon reacts with halogens to form tetrahalides ($SiX_4$).

  • Group 14 elements below carbon (like Silicon) cannot form strong $π$ bonds, leading to single-bond networks (e.g., $SiO_2$ vs. $CO_2$ molecules).

Compounds of Boron and Silicon

• Boron Halides: $BF_3, BCl_3, BBr_3, BI_3$. Trigononal planar, $sp^2$ hybridized. $BF_3$ is a Lewis acid that reacts with $F^-$ to form fluoroboric acid ($HBF_4$).

• Silicon Halides: $SiCl_4, SiF_4$. $SiCl_4$ is a low-boiling covalent liquid. $SiF_4$ is prepared via: $SiO_2(s) + 4HF(aq) \rightarrow SiF_4(g) + 2H_2O(l)$. This is why $HF$ cannot be stored in glass.

• Boron Oxides/Acids:

  • Boric oxide ($B_2O_3$) is used in heat-resistant borosilicate glass (Pyrex).

  • Boric acid ($B(OH)_3$) is a weak Lewis acid but not a proton donor in the classic sense: $B(OH)_3(aq) + 2H_2O(l) \rightleftharpoons B(OH)_4^-(aq) + H_3O^+(aq)$.

• Silicon Dioxide ($SiO_2$): Known as silica. Occurs as quartz (crystalline) or opal (amorphous). Forms a three-dimensional continuous network of $SiO_4$ tetrahedra.

• Silicates: Salts containing silicon and oxygen anions. Silicon is $sp^3$ hybridized at the center of tetrahedra.

Structure and Properties of Nonmetals

• Bonding Trends: Compound formed between metal and nonmetal is typically ionic (solid). Between two nonmetals, it is covalent (gas, liquid, or volatile solid).

• Ions: Nonmetals only form monatomic anions (e.g., $Cl^-, N^{3-}$), never monatomic cations.

• The "First Member" Rule: The first member of a group (Period 2) differs from others due to smaller size, higher ionization energy, and lack of $d$ orbitals (limiting coordination to 4).

• Allotropes of Carbon:

  • Diamond: Tetrahedral $sp^3$ network; very hard; melting point ~$4400 \text{ } ^\text{o} C$.

  • Graphite: Planar layers of six-membered rings; $sp^2$ hybridized; soft/flaky due to weak London forces between layers; conductive due to delocalized $π$ bonds.

  • Buckminsterfullerene ($C_{60}$): Soccer-ball-shaped icosahedral molecules.

  • Graphene: Single-atom thick layer of graphite; extremely strong and conductive.

  • Carbon Nanotubes: Graphene sheets rolled into tubes; walls can be one or multiple atoms thick.

Hydrogen: Preparation and Compounds

• Abundance: Most abundant element in the universe (90 \text{% } of atoms).

• Isotopes: Protium ($1H$), Deuterium ($2H$ or $D$), and Tritium ($3H$ or $T$).

• Industrial Preparation:

  • Water Gas: Passing steam over coke: $C(s) + H_2O(g) \rightarrow CO(g) + H_2(g)$.

  • Electrolysis: $2H_2O(l) + \text{electrical energy} \rightarrow 2H_2(g) + O_2(g)$.

  • Laboratory: Reaction of metals like iron with acid: $Fe(s) + 2H_3O^+(aq) + 2Cl^-(aq) \rightarrow Fe^{2+}(aq) + 2Cl^-(aq) + H_2(g) + 2H_2O(l)$.

• Hydrogenation: The addition of $H_2$ to reduce organic compounds.

• Hydrides:

  • Ionic Hydrides: Formed with Group 1 and heavier Group 2 metals ($H^-$ ion).

  • Ammonia ($NH_3$): Produced via Haber process: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$.

  • Hydrazine ($N_2H_4$): Colorless liquid used as rocket fuel.

  • Phosphine ($PH_3$): Toxic gas with a decaying fish odor.

  • Hydrogen Sulfide ($H_2S$): Rotten-egg odor; weak diprotic acid.

Nitrogen and Phosphorus Chemistry

• Nitrogen Fixation: Transformation of atmospheric $N_2$ into biologically useful chemicals ($NH_3$, nitrates) by bacteria (using nitrogenase enzymes) or lightning.

• Nitrogen Oxides:

  • $N_2O$ (Nitrous Oxide): "Laughing gas," dental anesthetic.

  • $NO$ (Nitric Oxide): Simplest stable molecule with an unpaired electron.

  • $NO_2$ (Nitrogen Dioxide): Brown, paramagnetic gas that exists in equilibrium with its dimer $N_2O_4$.

• Nitric Acid ($HNO_3$): Prepared via the Ostwald Process: $NH_3 \rightarrow NO \rightarrow NO_2 \rightarrow HNO_3$.

• Aqua Regia: Mixture of 1 part $HNO_3$ and 3 parts $HCl$; can dissolve gold ($Au$).

• Phosphorus Allotropes:

  • White Phosphorus ($P_4$): Tetrahedral molecules; waxy solid; waxy; bursts into flame in air; stored under water.

  • Red Phosphorus: Polymeric network of $P_4$ units; much less reactive and nontoxic.

• Phosphorus Oxides: $P_4O_6$ and $P_4O_{10}$. $P_4O_{10}$ is a powerful drying agent.

Oxygen and Sulfur Chemistry

• Oxygen: Most abundant element in the Earth's crust (50 \text{% } by mass).

• Ozone ($O_3$): Allotrope formed by electrical discharge in oxygen. Stratospheric ozone absorbs UV light. CFCs (Freons) deplete ozone by releasing chlorine atoms.

• Oxides, Peroxides, and Superoxides:

  • Oxide ($O^{2-}$): Standard binary compounds.

  • Peroxide ($O_2^{2-}$): Contains $O-O$ single bond (e.g., $Na_2O_2, H_2O_2$).

  • Superoxide ($O_2^-$): Bond order of $1.5$ (e.g., $KO_2$).

• Sulfur Allotropes: Stable form at room temperature is yellow rhombic sulfur ($S_8$ crown-shaped rings). Heating breaks rings into polymeric chains, increasing viscosity.

• Frasch Process: Mining of sulfur by forcing superheated water ($170 \text{ } ^\text{o} C$) and compressed air into underground deposits.

• Sulfuric Acid ($H_2SO_4$): Produced in higher quantities in industry than any other compound. Acts as a strong acid, dehydrating agent, and oxidizing agent.

Halogens and Noble Gases

• Halogens: $F_2$ (pale yellow gas), $Cl_2$ (greenish-yellow gas), $Br_2$ (reddish-brown liquid), $I_2$ (grayish-black solid). All are diatomic.

• Interhalogens: Compounds formed between different halogens ($XY_n$, where $n = 1, 3, 5, 7$). Iodine achieves the $7+$ state only with fluorine ($IF_7$).

• Noble Gases: Group 18. Helium, Neon, Argon, Krypton, Xenon, Radon.

• Compounds:

  • Bartlett proved noble gases could react in 1962.

  • Xenon forms fluorides ($XeF_2, XeF_4, XeF_6$) and oxides ($XeO_3$, which is explosive).

  • Krypton forms $KrF_2$.

• Applications:

  • Helium for balloons and cryogenic cooling for MRI magnets.

  • Neon for signs.

  • Argon for inert atmospheres in welding and light bulbs.