Pure Substances and Mixtures Quiz

Pure Substances and Mixtures

• A pure substance, scientifically, is composed of a single element or compound.
• Everyday usage of "pure" implies nothing is added, but scientifically, milk, for instance, contains a mixture.
• Pure substances have a sharp, precise melting point.
• Impure substances melt over a range of temperatures.

Identifying Pure Substances Using Melting Points

• A pure substance exhibits a definite, precise melting point.
• Mixtures melt across a range of temperatures, due to different components melting at various points.

Measuring Temperature

• Thermometers and temperature probes are used to measure temperature.
• Temperature probes give more precision, recording to 2 decimal places.

Relative Atomic Mass

• The average mass of an atom of an element compared to 1/12 the mass of a carbon-12 atom.

Relative Formula Mass

• The weighted average mass of the formula units in relation to 1/12 the mass of a carbon-12 atom.

Relative Molecular Mass

• The average mass of a molecule compared to 1/12 the mass of a carbon-12 atom.

Calculating Relative Formula Mass

• Add the relative atomic masses of each element in the formula.

Calculating the Relative Formula Mass of Ca(OH)2

• Ca(OH)2's relative formula mass is calculated as 40 + 2(16 + 1) = 74.

Empirical Formula

• The simplest whole-number ratio of atoms in a compound.

Molecular Formula

• The actual number of atoms of each element in a compound.

Empirical Formulae of CH4 and C4H10

• CH4: Already the simplest whole-number ratio, so the empirical formula is CH4
• C4H10: Simplifying the ratio gives C2H5

Molecular and Empirical Formulae of C2H4Br2

• Molecular formula: C2H4Br2
• Empirical formula: CH2Br

Alloys

• Mixtures of two or more metals.

Advantages of Alloys over Pure Metals

• Alloys often exhibit more desirable properties, such as increased hardness and strength.
• The different sizes of atoms in alloys distort the layers of the material, restricting movement and improving the material's strength.

Formulations of Mixtures

• Formulations are mixtures containing exact quantities of substances.
• These quantities are optimized to achieve the best properties for a specific purpose.

Examples of Formulations

• Sunscreen
• Medicine
• Perfume
• Drinks

Separating Insoluble Salts (Filtration)

• Filter paper in a funnel over a flask to separate solids from liquids.
• Pour the mixture through the funnel; wash remaining solids with distilled water to collect any remaining salt.
• Remove filter paper and evaporate water from the residue to obtain the salt.

Separating Soluble Salts (Crystallisation)

• Heat a solution gently to increase its concentration.
• Remove the solution from heat and allow it to cool slowly.
• Crystals will form as water evaporates.

Simple Distillation

• Used to separate miscible liquids with different boiling points.

Separating Ethanol from Water (Simple Distillation)

• Place mixture in a flask connected to a condenser with a collecting flask.
• Maintain a temperature gradient in the condenser by circulating cool water.
• Water will not evaporate due to its higher boiling point. Heat will evaporate the ethanol, which will then condense and collect.

Fractional Distillation (Crude Oil)

• Used to separate crude oil by boiling point differences.
• Vapours rise through a fractionating column with a temperature gradient.
• Different hydrocarbons condense at different fractions according to their boiling points.

Chromatography

• Separates mixtures of soluble substances by their different affinities for a stationary and mobile phase.

Paper Chromatography Phases

• Stationary phase: Paper
• Mobile phase: Solvent

Thin-Layer Chromatography (TLC) Phases

• Stationary phase: Thin layer of inert substance on a surface.
• Mobile phase: Solvent

Performing Paper/TLC Chromatography

• Draw a pencil line a few centimeters from the bottom of the stationary phase.
• Spot the mixture to be tested on the baseline.
• Place the stationary phase in a beaker with solvent, ensuring the solvent level is below the pencil line.
• The solvent will travel up the paper, separating substances.

Using Pencil in Chromatography

• Pencil is insoluble and won't affect the chromatography's results.
• Ink is soluble and can interfere with the results.

Rf Value

• Ratio of distance travelled by solute to the distance travelled by the solvent.

Calculating the Rf Value

• Rf = Distance travelled by substance / Distance travelled by solvent

Gas Chromatography (GC) Use

• Separates mixtures of volatile liquids.

Gas Chromatography Phases

• Stationary phase: Solid/liquid on a solid support.
• Mobile phase: Inert carrier gas.

Gas Chromatography Separation

• Substances travel through a column at different speeds.
• Retention time (time to reach the detector) identifies them.

Using Chromatography to Determine Purity

• Pure substances show one spot/peak.
• Impure substances show multiple spots/peaks.

Metals in the Periodic Table

• Found on the left side of the periodic table.
• React to form positive ions.

Non-Metals in the Periodic Table

• Found on the top right of the periodic table.
• React to form negative ions.

General properties of Metals

• Shiny
• Good conductors
• Dense
• Malleable and ductile
• High melting and boiling points

General properties of Non-metals

• Dull appearance
• Poor conductors
• Lower density than metals
• Low melting and boiling points
• Brittle

Metal Reaction with Oxygen

• Metal oxide is formed.

Ion Formation

• Positive ions (cations) are formed when a metal loses electrons.
• Negative ions (anions) are formed when a non-metal gains electrons.

Magnesium and Oxygen Reaction

• 2Mg + O2 → 2MgO

Periodic Table Arrangement of Elements

• Ordered by increasing atomic number.
• Elements in the same group have similar properties.

Similarity of Properties in the Same Group

• Elements have the same number of outer shell electrons, influencing reactions.

Periodic Table Period/Row

• Period number indicates the number of electron shells possessed by the elements in that row.

Periodic Table Group/Column

• Group number indicates the number of outer shell electrons.

Covalent vs. Ionic Bonds

• Covalent: Non-metals share electrons.
• Ionic: Metal loses electrons to non-metal creating oppositely charged ions.

Ionic Compound Bonding

• Electrostatic attraction between positive and negative ions.

Ionic Compound Properties (High Melting/Boiling Points)

• Strong electrostatic forces require significant energy to overcome during melting/boiling.

Ionic Compound Conductivity

• Conduct electricity when molten or dissolved (aqueous), as the ions are free to move and carry charge.

Bonding in Simple Molecules

• Covalent bonds form from electron sharing between non-metals.

Simple Molecules (Low Boiling Points)

• Weak intermolecular forces between molecules,requiring less energy for boiling.
• The covalent bonds within the molecules remain intact.

Simple Molecules and Electricity Conductance

• Simple molecules do not conduct electricity as they lack an overall charge.

Simple Molecule Boiling Point Change with Size

• Stronger intermolecular forces lead to higher boiling points with increasing molecular size.

Giant Covalent Structures (Bonding)

• Many strong covalent bonds connecting atoms in a large network.

Giant Covalent Structures (Melting Points)

• High melting points because breaking covalent bonds requires lots of energy.

Polymer Bonding

• Covalent bonds within the polymer's structure.

Polymer Solids

• Strong intermolecular forces, holding large polymer molecules together in a solid form.

Metallic Bonding

• Metal atoms form a giant structure of positive ions in a sea of delocalized electrons.

Metallic Conductivity

• Delocalized electrons moving through the structure carry charge for conductivity.

Metallic Malleability

• Layers of atoms in uniform arrangement can slide easily past each other.

Metallic Melting Points

• Strong metallic bonds, resulting in relatively high melting points requiring large amounts of energy to overcome the attraction.

Dot and Cross Diagrams Limitation

• Do not show the 3D arrangement of atoms within the molecule.

Ball and Stick Models

• Show the 3D shape of the molecule and the bonds between atoms, but do not represent electrons.

Highest Electron Configuration (Three Shells)

• 2, 8, 8.

Most Desirable Electron Configuration

• Full outer shell, typically 8 electrons.

Noble Gas Reactivity

• Stable full outer shell configurations make them unreactive.

Reactive Element (Example)

• An element with 2, 8, 1 electrons configuration can easily lose one electron to attain a stable configuration.

Mendeleev's Periodic Table

• Ordered elements by increasing atomic mass, adjusting to place elements with similar properties in the same group.
• Left gaps for undiscovered elements predicting their properties.

Modern Periodic Table vs. Mendeleev's Table

• Modern table is organized by increasing atomic number.

Carbon Bonding

• Carbon can form four covalent bonds.

Organic Compounds

• Compounds containing carbon covalently bonded to other atoms.

Diverse Organic Compounds

• Carbon's ability to form chains, rings, and families of similar compounds results in diverse organic compounds.

Graphite Structure

• Carbon atoms arranged in hexagonal layers with one delocalized electron per carbon atom.

Graphite Properties

• Soft/slippery layers allow for sliding, conducting electricity due to delocalized electrons.

Diamond Structure

• Each carbon atom is bonded to four other carbon atoms, forming a giant tetrahedral structure with no delocalized electrons.

Diamond Properties

• Very hard and high melting point due to strong covalent bonds.
• Does not conduct electricity.

Fullerene

• Molecule of carbon with a closed tube or hollow ball shape.

Examples of Fullerenes

• Graphene, C60 (buckminsterfullerene).

Fullerene Properties and Uses

• Large surface area for catalyst trapping.
• Hollow structure for targeted drug delivery.

Graphene in Electronics

• Strong, conductive, and only one atom thick, making it useful in electronics.

Energy Transfer in State Changes

• Energy is transferred from the surroundings during condensation and freezing.
• Energy is transferred to the substance during evaporation and melting.

Substance State at a Temperature

• Substance A (-174°C): Liquid
• Substance B (-7°C): Solid

Atom vs. Substance Properties

• Individual atoms don't have the same characteristics as the material they form. Physical properties depend on bonding and the structure of the substance.

Nanoparticle Size

• 1–100 nanometers across, larger than typical atoms but smaller than most particles.

Surface Area to Volume Ratio

• Nanoparticles have a much higher surface area to volume ratio compared to larger particles.

Nanoparticle Reactivity

• High surface area to volume ratio leads to a higher number of accessible reaction sites, increasing reactivity.

Nanoparticles as Catalysts

• High surface area to volume ratio maximizes reaction sites for catalysis.

Surface Area to Volume Ratio Calculation

• Surface area to volume ratio = surface area ÷ volume.

Nanotubes in Electrical Circuits

• Lightweight, conductive, and small enough for use in computer circuits.

Nanoparticles in Sunscreen

• Use high surface area to volume ratio to block UV light without creating white marks..

Nanoparticle Risks

• Limited understanding of their long-term risks.
• Potential for harmful effects if they are able to enter the bloodstream.