Metallic Bonding and Properties of Metals

Metallic Bonding

• Metals have atoms closely packed in layers, creating a metallic lattice.
• Atoms lose electrons from outer shells, forming positive metal ions.
• Metallic bonding comes from the electrostatic attraction between positive ions and delocalized electrons.
• Delocalized electrons move freely in the metal structure.
• Electrical Conductivity: Ability to conduct an electric current.
• Delocalized Electrons: Electrons free to move throughout the structure.
• Ductility: Property of metal allowing it to be drawn into wires.
• Malleability: Property of metal allowing it to be hammered into shapes.
• Metallic Bonding: Electrostatic forces between positive ions and delocalized electrons.

Properties of Metals

• Delocalized electrons allow metals to conduct electricity.
• Metals are malleable, bending without breaking due to sliding ion layers and maintaining structure.
• Metals are ductile, able to be drawn into thin wires for electrical wiring.
• Metals have high density due to tight packing and are shiny when the oxide layer is removed.
• Transition metals are denser than water and sink.
• Metals have high melting and boiling points because of strong electrostatic forces.
• Mercury is liquid at room temperature.

Factors Affecting Melting Points

• More delocalized electrons lead to higher melting points.
• Higher metal cation charges result in stronger electrostatic forces and higher melting points.

Reactivity Series

• The reactivity series is a list of metals from most to least reactive.
• Potassium is most reactive, gold is least reactive.
• Reactivity is determined by observing reactions with water, acids, and oxygen.
• Potassium, sodium, and lithium react violently with cold water, producing metal hydroxide and hydrogen gas.
• Magnesium reacts with steam to form magnesium oxide and hydrogen gas.
• Magnesium and aluminum react slowly or not at all with water due to an oxide layer.
• Copper, silver, and gold don't react with dilute acids.
• Sodium and potassium react violently with dilute acids.
• Most metals react with dilute acids to produce salt and hydrogen gas.
• Reacting various metals with acids indicates the reactivity order, with more vigorous reactions indicating more reactive metals.

Displacement Reactions

• Displacement reactions involve reacting a more reactive metal with the metal oxide of a less reactive metal, such as aluminum displacing iron from iron oxide.
• Magnesium is more reactive than copper.
• Copper loses oxygen and is reduced in the reaction.
• Magnesium gains oxygen and is oxidized.
• Oxidation and reduction result in a redox reaction.

Oxidizing and Reducing Agents

• An oxidizing agent gives oxygen to another substance and becomes reduced.
• A reducing agent accepts oxygen and becomes oxidized.
• OILRIG explains oxidation and reduction: Oxidation Is Loss, Reduction Is Gain (of electrons).

Displacement Reactions with Salt Solutions

• Reactions involve a more reactive metal displacing a less reactive metal from its salt solution.
• Spectator ions are unchanged in a chemical reaction.

Metal Reactivity, Rusting, and Extraction

• Metals are arranged in a reactivity series based on displacement reactions with metal salts (Potassium > Sodium > Lithium > Calcium > Magnesium > Aluminum > Zinc > Iron > Copper > Silver > Gold).
• Metals can be arranged in a reactivity series based on their displacement reactions with aqueous solutions of metal salts.
• Redox (oxidation-reduction) reactions involve the transfer of electrons.
• Oxidation: Loss of electrons.
• Reduction: Gain of electrons.
• Oxidizing Agent: A substance that causes oxidation by accepting electrons.
• Reducing Agent: A substance that causes reduction by donating electrons.
• Corrosion is oxidation and weakening over time.
• Rusting refers to the corrosion of iron or steel when exposed to oxygen and water.
• Iron is oxidized to form Fe3+ ions, reacting with water to form hydrated iron(III) oxide (rust)
• Chemical formula for rust is Fe2O3·xH2O.

Preventing Rusting

• Barrier methods: Coating with oil, paint, grease, or plastic to prevent water and oxygen.
• Galvanizing: Coating iron or steel with zinc, which acts as a barrier and a sacrificial metal.
• Sacrificial protection: Attaching a more reactive metal to iron or steel, which oxidizes more easily and protects the iron.

Metal Extraction

• Most metals are extracted from ores, while unreactive metals are found as native metals.
• Extraction Method depends on the metal's position in the reactivity series:
• Carbon extraction is used for metals below carbon and involves heating the ore with carbon.
• Electrolysis is used for metals higher than carbon and requires electrolysis.
• Electrolysis is a chemical change caused by passing an electric current through a molten or dissolved compound.
• This is essential for extracting metals more reactive than carbon, such as aluminum.

Uses of Metals and Alloys

• An alloy is a mixture of a metal and one or more other elements, usually other metals or carbon.
• Alloys are harder than pure metals due to disrupted atom arrangements.
• Steel is an alloy containing about 4% carbon, phosphorus, and other elements.
• Adding other elements to iron creates steels with varied properties.
• Stainless steel contains chromium, which forms a chromium oxide layer (Cr2O3) that prevents rusting.

Properties and Uses of Common Metals

• Aluminum: low density, good conductor, corrosion-resistant (Electrical cables, pots and pans, aluminum foil, aircraft)
• Copper: antimicrobial, very good conductor of heat and electricity, ductile (Electrical wires, pots and pans, water pipes, surfaces in hospitals)
• Iron: malleable, strong, brittle (Fences, making steel, construction)

Identifying Unknown Substances

• Chemists identify substances using various tests in industries.

Tests for Gases

• Hydrogen: Place a lit splint at the mouth of the boiling tube; squeaky pop sound indicates hydrogen.
• Oxygen: Place a glowing splint into the boiling tube; glowing splint relights indicate oxygen.
• Carbon Dioxide: Bubble the gas through limewater; colorless and clear limewater turns cloudy white.
• Ammonia: Place damp red litmus paper at the mouth of the boiling tube; paper turns blue.
• Chlorine: Place damp blue litmus paper at the mouth of the boiling tube; paper turns red then white (bleaching effect).

Tests for Positive Ions

• Identifying salts can be challenging because many salts are white crystalline solids.

Key Terminology

• Nichrome is an unreactive metal alloy.
• A precipitate is a fine insoluble solid formed by a chemical reaction in solution.
• Cations can be identified using a Bunsen burner flame, aka the flame test.
• Procedure for Flame Test:
  • Clean a platinum or nichrome wire in hydrochloric acid.
  • Moisten the wire with hydrochloric acid and dip it in the test salt you want to test.
  • Place the wire with the salt test in a blue Bunsen burner flame, and observe the color of the flame.
  • Flame colors are:
    • Li+ = Red
    • Na+ = Yellow
    • K+ = Lilac
    • Ca2+ = Orange-Red
    • Cu2+ = Blue-Green

Using Sodium Hydroxide to Identify Cations

• Add unknown salt in water, add 1 cm3 of the salt solution to a test tube, add 1cm3 of sodium hydroxide solution to the test tube, and observe precipitate color formed.

Identifying Ammonium Ions Using Sodium Hydroxide

• Add 2cm of unknown solution to test tube; add drops of sodium hydroxide solution, shake well, warm with flame, hold damp red litmus paper near the mouth of the test tube; paper turns blue due to ammonia gas produced in the test.
• Sodium hydroxide solution cannot be used to identify Li+ and K+ ions, but flame tests can.