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Questions and Answers
What feature contributes to the low melting and boiling points of simple molecular substances?
What feature contributes to the low melting and boiling points of simple molecular substances?
Which of the following is an example of an allotrope of carbon?
Which of the following is an example of an allotrope of carbon?
What structure characterizes Buckminsterfullerene?
What structure characterizes Buckminsterfullerene?
Which property distinguishes ionic substances from simple molecular solids?
Which property distinguishes ionic substances from simple molecular solids?
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What type of structure do ionic compounds typically form?
What type of structure do ionic compounds typically form?
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Which of the following describes a property of graphite?
Which of the following describes a property of graphite?
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Which statement is true regarding the allotropes of sulfur?
Which statement is true regarding the allotropes of sulfur?
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Which of the following features is NOT characteristic of simple molecular solids?
Which of the following features is NOT characteristic of simple molecular solids?
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What is the primary reason that simple molecular crystals have lower melting and boiling points compared to giant covalent solids?
What is the primary reason that simple molecular crystals have lower melting and boiling points compared to giant covalent solids?
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Which of the following substances is an example of a simple molecular crystal?
Which of the following substances is an example of a simple molecular crystal?
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What type of intermolecular forces are present between the molecules in simple molecular crystals?
What type of intermolecular forces are present between the molecules in simple molecular crystals?
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At what temperature does white phosphorus (P4) melt?
At what temperature does white phosphorus (P4) melt?
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Which of the following statements correctly describes iodine (I2) as a simple molecular solid?
Which of the following statements correctly describes iodine (I2) as a simple molecular solid?
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What distinguishes simple molecular crystals from giant covalent solids?
What distinguishes simple molecular crystals from giant covalent solids?
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Which of the following is NOT a characteristic of simple molecular crystals?
Which of the following is NOT a characteristic of simple molecular crystals?
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Which of the following correctly describes the molecular structure of sulfur in solid form?
Which of the following correctly describes the molecular structure of sulfur in solid form?
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What describes the arrangement of Na+ and Cl- ions in the crystalline structure of sodium chloride?
What describes the arrangement of Na+ and Cl- ions in the crystalline structure of sodium chloride?
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Which of the following accurately describes the type of bonding in sodium chloride?
Which of the following accurately describes the type of bonding in sodium chloride?
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Which characteristic is primarily responsible for the high melting and boiling points of sodium chloride?
Which characteristic is primarily responsible for the high melting and boiling points of sodium chloride?
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What process occurs when sodium chloride dissolves in water?
What process occurs when sodium chloride dissolves in water?
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In terms of intermolecular forces, which of the following is strongest?
In terms of intermolecular forces, which of the following is strongest?
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What defines the unit cell of a crystal such as sodium chloride?
What defines the unit cell of a crystal such as sodium chloride?
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Which of the following statements about the structure of NaCl is incorrect?
Which of the following statements about the structure of NaCl is incorrect?
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Which factor is least relevant to the differentiating properties of molecular structures?
Which factor is least relevant to the differentiating properties of molecular structures?
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Study Notes
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?
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