Untitled Quiz

Crystal Structures of Substances

• Ionic solids have a crystal lattice structure
• Allotropes are different structural forms of the same element in the same physical state.
• Simple molecular solids consist of molecules held together by weak intermolecular forces
• Giant covalent solids are composed of atoms bonded together by strong covalent bonds in a regular 3-dimensional arrangement.
• Amorphous solids lack a crystalline structure.

Objectives

• Describe the structure of ionic crystals and give examples
• Relate the properties and use of sodium chloride to its structure
• Make diagrammatic representations of sodium chloride
• Make models of sodium chloride in groups for grading
• Describe the structure of simple molecular crystals and give examples of each
• Distinguish between ionic and simple molecular crystals by use of melting point, solubility, and conductivity
• Investigate melting point, solubility of solids, and conductivity of resulting solutions
• Describe the structure of giant molecular crystals and give examples
• Explain allotropy and give examples of allotropes of carbon (diamond and graphite)

Ionic Crystals/Ionic Substances

• Formed by ionic bonding
• Composed of an ionic lattice
• Cations and anions arranged in a repeating three-dimensional pattern
• Held together by strong electrostatic forces (ionic bonds)
• Represented by empirical formulae or formula units.
• Examples include sodium chloride.

Structure and Properties of Ionic Substances (eg. sodium chloride)

• Sodium chloride is a typical ionic compound
• Ions in sodium chloride are held in a regular lattice
• The ions are held by strong ionic bonds
• The lattice grows in all directions, giving a crystal
• Crystals are white and shiny
• Uses of NaCl include use as an electrolyte such as in water

Properties and uses of Ionic Compounds

• High melting point due to strong electrostatic forces
• Hard and brittle.
• Crystalline solids.
• Soluble in water (polar).
• Conduct electricity when molten or dissolved in water.

Dissolution and Solvation of NaCl

• Solvation is the process of surrounding solute particles with solvents
• Solvation in water is called hydration.
• Attraction between water dipoles and the ions of a crystal is greater than the attraction among the ions of a crystal.

Lattice Structure of NaCl

• The given figure clearly illustrates the unit cell structure of NaCl
• There are four NaCl molecules per unit cell
• Each Na+ ion is surrounded by 6 Cl- ions
• Each Cl- ion is surrounded by 6 Na+ ions

Model of Sodium Chloride Structure

• Students can use materials like matches sticks, icicle pops, flour dough or play dough.
• Model should include correct amount and positions of Na+ and Cl- ions
• Due date Friday 17th

Intramolecular vs Intermolecular Forces

• Intramolecular forces hold atoms within a molecule
• Intermolecular forces hold molecules together
• Intermolecular forces are weaker compared to intramolecular forces
• Forces are categorized into dipole-dipole, London dispersion, and hydrogen bonding forces

Weak Intermolecular forces

• Arise from attraction between dipoles in neighboring molecules.
• Types include hydrogen bonding, permanent dipole-dipole, van der Waals forces

Strength Differences of Intermolecular Forces

• Intermolecular forces are weaker compared to covalent, ionic and metallic forces
• Strength order is: hydrogen bonding > permanent dipole-dipole > van der Waals

Examples of Intermolecular Forces

• Hydrogen chloride (HCl), hydrogen fluoride (HF), and water (H₂O) are all examples of dipole-dipole forces.

Van der Waals Forces

• Not permanent
• Electrons in atoms are in constant motion
• Instantaneous dipole is formed.
• The dipole can induce the formation of a dipole in neighboring molecules.
• Neighboring molecules attract each other because of their dipoles.
• London dispersion forces are a type of van der Waals force.

London Dispersion Forces

• Temporary attractive force between adjacent atoms
• Occurs when electrons in adjacent atoms occupy positions creating temporary dipoles
• Results in induced dipole-induced dipole attraction
• Weakest of intermolecular forces

Useful Information

• Water's unusual properties (e.g., lower ice density and high surface tension) are due to extensive hydrogen bonding

Hydrogen Bonding

• Special form of permanent dipole bonding
• Requires one molecule with an H atom covalently bonded to an F, O, or N atom
• Also requires a second molecule with a F, O, or N atom and a lone pair of electrons.

Hydrogen Bonding in Water

• Hydrogen bonding in water affects its boiling point
• Boiling point is high due to the strong hydrogen bonds

Simple Molecular Crystals

• Covalently bonded substances, can be simple molecular crystals or giant covalent solids.
• Simple molecules are held together by weak intermolecular forces.
• Examples include carbon dioxide, water, sulfur, phosphorus, iodine and short chain hydrocarbons.

Example of simple molecular crystals

• Water (ice)
• Solid carbon dioxide (dry ice)
• Sulfur (S8)
• Phosphorus (P4)
• Iodine (I2)
• Sucrose (table sugar)
• Short-chain hydrocarbons (e.g., methane)

Structure of White Phosphorus

• Consists of tetrahedral P4 molecules
• Molecules held together by weak van der Waals forces
• Melting point: 44.1 °C, Boiling point: ~280 °C

Structure of Iodine

• Crystalline solid with a purple vapor
• Melting point: 114°C, Boiling point: 184°C
• Slightly soluble in water but dissolves in organic solvents.

Question about Simple Molecular Crystals

• Why do simple molecular crystals have low melting and boiling points?
• What are some examples of simple molecular crystals?

Differentiating between Ionic and Covalent Substances

• Ionic substances have crystal lattices; covalent substances have molecules.
• Lithium chloride is an example of ionic; water is an example of covalent

Properties of Simple Molecular Solids

• Type of Chemical Bond: Covalent bonds hold atoms in molecules; weak intermolecular forces hold molecules together
• Melting Point: Low, due to weak intermolecular forces
• Solubility: Polar substances dissolve in polar solvents, and non-polar substances dissolve in non-polar solvents
• Conductivity: Do not conduct electricity in any state because they contain no charged particles (free ions/electrons)
• State: Can exist as soft solids, gases or liquids.

Summary: Ionic vs. Simple Covalent Substances

• State at room temperature (Ionic = Crystalline solids; Covalent = Gases, liquids, or low melting solids)
• Polarity (Ionic=Polar; Covalent = Polar or non-polar)
• Solubility (Ionic = Frequently soluble in water and organic liquids; Covalent = Few are soluble in water)
• Conductivity (Ionic = Good conductors in molten state or dissolved in water; Covalent = Do not conduct electricity)
• Melting point (Ionic = High; Covalent = Low)
• Boiling point (Ionic = High; Covalent = Low)
• Examples (Ionic = NaCl; Covalent = CH4)

Giant Atomic Crystals

• Composed of non-metal atoms bonded together by strong covalent bonds
• Also known as macromolecule
• Examples include diamond, graphite, and silicon dioxide (silica)

Silicon Dioxide (Silica/Quartz)

• Abundant Earth mineral
• Main component of rocks and sand
• Forms crystalline quartz.

Properties of Diamond

• Very high melting point (3550°C)
• Extremely hard
• Does not conduct electricity (Valence electrons are bound)

Properties of Graphite

• Very high melting point (3600°C)
• Soft and lubricating properties (Layers slide)
• Good conductor of electricity (De-localized electrons)

Allotropy

• Property of some elements to exist in two or more different forms (allotropes) in the same physical state.
• Examples include allotropes of carbon (diamond, graphite, fullerenes), sulfur (rhombic and monoclinic), and phosphorus (white and red).

Allotropes of Carbon

• Diamond, graphite, and fullerenes are allotropes of carbon

Allotropes of Sulphur

• Rhombic & monoclinic are allotropes of sulfur.

Allotropes of Phosphorus

• White phosphorus and red phosphorus are allotropes of phosphorus.

Giant Metallic Crystal Lattices and Characteristics

• Metal atoms are arranged in compact, orderly patterns
• Held together by electrostatic forces between metal cations and negatively charged electrons

Characteristics of Metals

• Solids at room temperature (except mercury)
• Malleable (can be hammered into thin sheets)
• Ductile (can be drawn into wires)
• High melting and boiling points
• High density
• Good conductors of heat and electricity

Malleability of Metals

• Metal atoms move to new locations, allowing for the change in the shape of a metal.

High Melting Points of Metals

• High melting points indicate strong metallic bonding.
• Melting points increase with the number of valence electrons in the 'sea' of electrons.

High Density of Metals

• Metal lattices are close-packed, contributing to high densities.

Conductivity (Heat and Electricity) in Metals

• Free-moving electrons in metals allow for efficient heat and electrical conductivity.

Alloys

• Alloys are mixtures of two or more elements, at least one of which is a metal.
• They are usually stronger than their individual components.
• Examples are steel (iron, carbon, other elements)

Summary Questions

• What type of crystal structure does each of the following have? (Potassium bromide, Iodine, Ice, Graphite, Calcium carbonate)
• Why can sodium chloride dissolve in water?
• Compare the melting points of an ionic solid and a simple molecular solid.
• Why are diamonds extremely hard?
• Why can graphite conduct an electric current?