Chemistry: Pure Substances and Mixtures

• Chemistry notes for test prep excluding hydrocarbons

Pure Substances

• Made of only one type of particle
• Have a fixed composition and distinct physical/chemical properties

Types of Pure Substances

• Elements: One type of atom (e.g., O₂, Au, C). Cannot be broken down by chemical means
• Compounds: Two or more elements chemically bonded (e.g., H₂O, CO₂, NaCl). Can be broken into elements via chemical reactions.

Mixtures

• Combination of two or more substances physically mixed
• Components retain their individual properties
• Can be separated by physical means

Types of Mixtures

• Homogenous (Solutions): Uniform composition, components not visibly distinct (e.g., saltwater)
• Heterogeneous: Non-uniform, components are visibly distinct (e.g., sand and salt)

Properties of Pure Substances

• Melting/Boiling Points: Unique and constant for each pure substance under standard conditions; Used to identify substances
• Density: Constant ratio of mass to volume; Density = Mass / Volume
• Reactivity: Predictable chemical reactions (e.g., Mg reacts with O₂ to form MgO).
• Strength: Consistent resistance to stress or deformation (e.g., Titanium is strong and used in aerospace)
• Solubility: Specific ability to dissolve in a solvent (e.g., Sugar dissolves in water)
• Electrical Conductivity: Some pure substances will conduct electricity (e.g., Cu), some won't (e.g., SiO₂)
• Magnetism: Some are magnetic (e.g., Fe), it helps with identification;
• Optical Properties: Light interaction (e.g., color, transparency). Quartz is transparent with a specific refractive index

Properties of Mixtures

• Lack fixed properties, depends on components
• Can be separated by physical techniques (e.g., distillation, filtration)
• Salt increases the boiling point of water; concrete strength varies with water-cement ratio

Separation Techniques for Mixtures

• Magnetic Separation: Magnetism is leveraged (e.g., iron filings from sand)
• Sieving: Particle size difference is leveraged (e.g., gravel from sand)
• Decantation: Density difference is leveraged (e.g., sand and water)
• Filtration: Particle size is leveraged (e.g., sand from water)
• Evaporation/Crystallization: Volatility is leveraged (e.g., salt from saltwater)
• Distillation: Boiling point differences are leveraged (e.g., ethanol from water)
• Precipitation: Solubility limits are leveraged (e.g., BaSO₄ from solution)
• Centrifugation: Density is leveraged (e.g., blood cells from plasma)
• Chromatography: Solubility and attraction is leveraged (e.g., ink pigments)
• Solvent Extraction: Solubility in different solvents is leveraged (e.g., caffeine from coffee)

Types of Bonding

• Ionic: Electron transfer (metal + non-metal)
• Covalent: Electron sharing (non-metals)
• Metallic: Delocalized electrons in metals

Valency

• Bonding capacity of an atom
• H = 1, O = 2, N = 3, C = 4

Electronegativity and Bonding

• Ionic: ΔEN > 1.7 (e.g., NaCl)
• Polar Covalent: 0.5 ≤ ΔEN ≤ 1.7 (e.g., H₂O)
• Non-Polar Covalent: ΔEN < 0.5 (e.g., O₂)

Metallic Bonding and Properties

• Structure: Lattice of metal ions + sea of delocalized electrons
• Properties: Malleability and Ductility, Thermal/Electrical Conductivity, Luster, High Melting/Boiling Points

Alloys

• Mixture of metals or metal + non-metal
• Purpose: Enhanced strength, durability, corrosion resistance
• Types: Substitutional (e.g., Brass = Cu + Zn), Interstitial (e.g., Steel = Fe + C)

Covalent Substances

• Molecular: Discrete molecules, low melting/boiling points, poor conductivity (e.g., H₂O, CO₂)
• Network: 3D bonded lattice, high melting points, hard, poor conductivity (e.g., Diamond, SiO₂)

Allotropes of Carbon

• Diamond: Hard, non-conductive
• Graphite: Soft, conductive
• Fullerenes: Moderate hardness, conductive

Ions

• Cations: Positive ions (metal loses electrons)
• Anions: Negative ions (non-metal gains electrons)
• Polyatomic Ions: Group of atoms with net charge (e.g., SO₄²⁻)
• Neutral overall, charges balanced
• Properties: High melting points, brittle, conductive in molten/solution state

Naming Compounds

• Covalent: Prefixes (mono-, di-, tri-, etc.) are used; the Second element ends in "-ide"; CO is carbon monoxide and SF₆ is sulfur hexafluoride
• Ionic: Metal + Non-metal ("-ide"); Transition metals use Roman numerals; Polyatomic ions retain their name

Writing Formulas

• Balance charges using subscripts
• Use brackets for multiple polyatomic ions
• Examples: Al₂O₃, FeCl₃, Ca(NO₃)₂

Melting and Boiling Points

• Melting Point (MP): Temperature at which a substance changes from solid to liquid
• Boiling Point (BP): Temperature at which a substance changes from liquid to gas
• Pure substances have sharp and fixed MPs/BPs under standard pressure (1 atm); Used as identifiers (e.g., Ethanol BP = 78.5°C)
• Mixtures don't have fixed MP or BP and vary depending on composition (e.g., Saltwater has higher BP than pure water)
• Bond Strength: Stronger bonds will result in higher MP/BP
• Ionic and covalent networks have high MPs/BPs
• Molecular substances have low MPs/BPs due to weak forces
• Intermolecular Forces: Stronger forces result in a Higher MP/BP
• Impurities: Lower MP, broaden melting range (used in purity testing)
• Applications: Used in material selection; separation technique (e.g., distillation)

Bond Type and MP/BP

• Ionic: High melting/boiling point
• Covalent (Molecules): Low melting/boiling point
• Covalent (Networks): Very High melting/boiling point
• Metallic: Generally High melting/boiling point

Applications Based on Properties

• Ionic Compounds: Salts, electrolytes
• Covalent Molecules: Medicines, plastics
• Covalent Networks: Tools (e.g., diamond)
• Metals/Alloys: Construction, wiring, cookware

Definitions of Melting and Boiling

• Melting Point: Temperature at which a solid turns into a liquid; Particles have enough energy to overcome attractive forces holding them in a fixed position
• Boiling Point: Temperature at which a liquid turns into a gas; Particles gain enough energy to break free from the liquid's surface and enter the gas phase

Factors Affecting Melting & Boiling Points

• Melting and boiling points are physical properties of a substance
• Used to identify pure substances because they occur at fixed, characteristic temperatures under standard conditions (1 atm pressure)
• Mixtures have variable melting/boiling ranges due to the different substances involved
• Strong electrostatic forces exist between oppositely charged ions in Ionic Bonds
• Weak intermolecular forces (e.g., Van der Waals, hydrogen bonding) exist between covalent molecules
• Very High Melting/Boiling points are caused by extensive strong covalent bonds in 3D network structure
• Metallic Bonding: strong attraction exists between cations and delocalized electrons
• Stronger intermolecular forces result in a higher melting/boiling point
• Water (H₂O) has hydrogen bonding making it have a higher boiling point than CH₄; Larger molecules have more London dispersion forces, this means they have have higher melting/boiling points
• Larger molecules have more electrons, stronger London dispersion forces, and higher melting/boiling points
• Increasing pressure raises the boiling point
• Boiling occurs when vapor pressure = atmospheric pressure

Trends in the Periodic Table

• Metals (Group 1): Lower melting/boiling points down the group
• Non-metals (Group 17): Boiling point increases down the group. Larger molecules have stronger dispersion forces
• Melting/boiling points vary across a period depending on bonding type
• Middle of the period (metals) have the highest melting/boiling points.
• Right side (non-metals) have lower points due to covalent molecular bonding

Special Cases

• Water (H₂O) has a Boiling Point of 100°C and Melting Point of 0°C; Exceptionally high due to hydrogen bonding
• Mercury (Hg) has a MP of -38.8°C and a BP of 357°C; Only metal liquid at room temperature

Exam Tips

• Pure substances melt/boil at fixed temperatures; Mixtures have a range
• Compare bond types first to predict which has the higher melting/boiling point
• Use molecular size and intermolecular forces for covalent substances
• Use known examples (NaCl, H₂O, O₂, Diamond, Iron)

Summary Points

• Melting/boiling points reflect the energy needed to overcome forces holding particles together
• Ionic, metallic, and covalent network substances have high melting/boiling points due to strong bonds
• Covalent molecules have low melting/boiling points due to weak intermolecular forces
• Knowledge of bonding types helps predict material properties