Chemistry Module 1: Properties and Structure of Matter

Mixtures and Separation Techniques

Mixtures and Pure Substances

• Mixture: Substance that can be physically separated into two or more different substances (e.g., salt water).

• Pure substance: Cannot be separated into simpler substances by physical techniques; made up of the same atom or same unit throughout (e.g., water).

Homogeneous and Heterogeneous Mixtures

• Homogeneous mixtures: Components are uniformly distributed throughout (e.g., alloys, solutions, air).

• Heterogeneous mixtures: Components of the mixture are not uniform or have localized regions with different properties (e.g., cereal, sandstone).

• Suspension: A type of heterogeneous mixture, consisting of insoluble solid particles in a liquid.

Common Separation Techniques

Separation techniques exploit differences in physical properties like size, density, solubility, and magnetism.

• Sieving: Separates substances with different particle sizes (e.g., flour from wheat).

• Decanting: Separates a suspension where denser solid particles settle, allowing the liquid to be poured off (e.g., separating wine from sediment). Decanting and separating funnels rely on settling and density.

• Separating Funnel: Used for immiscible liquids with unequal densities; the denser liquid forms the bottom layer (e.g., oil and water separation).

• Magnetic Separation: Separates mixtures of two solids where one component has magnetic properties (e.g., iron filings from sand). This technique relies on magnetic properties.

• Filtration: Used to separate insoluble solids from liquids or soluble substances from insoluble ones. This technique relies on insolubility and particle size.

  • Example: Mixture of sand and water. Filtration is used in water treatment plants to remove solid particles.

  • Tip: Rinse residue with minimal, cold solvent; weigh filter paper before use.

Advanced Separation Techniques

• Evaporation: Separates a soluble solid from a solvent to obtain the solute (e.g., salt from saltwater solution).

  • Tip: Do not evaporate to dryness when collecting the solid; weigh the evaporating basin before use.

• Distillation: Separates liquids from solids in solution or two or more liquids from one another based on differences in boiling points (e.g., salt and water). This is effective for miscible liquids with large differences in boiling points.

• Fractional Distillation: A refined distillation technique used to separate liquids with very close boiling points. It involves multiple distillation steps (e.g., separation of crude oil, ethanol, or nitrogen from carbon dioxide). Distillation relies on boiling point differences, and fractional distillation uses a column to improve separation when boiling points are close.

Chemical Composition and Nomenclature

Percentage Composition

• Calculating percentage composition by weight: \text{Percentage by weight} = \frac{m{\text{component}}}{m{\text{sample}}} \times 100 \text{%}

• This calculation is a key part of Gravimetric analysis.

Inorganic Compounds

• Typically chemical compounds that lack carbon-hydrogen bonds; they can be ionic or covalent.

Naming Compounds

• Ionic Compounds:

  • Name the positive ion first, then the negative ion (first syllable) with an "-ide" suffix (e.g., NaCl \rightarrow Sodium Chloride).

  • Ignore the number of atoms in the name.

  • If the metal can have more than one valency, the valency is shown in brackets (e.g., Iron(II) sulfide). Ionic naming uses valency in brackets when variable.

• Covalent Compounds:

  • Elements are named in order of least to most electronegative.

  • Naming rule: (prefix if > 1 atom)(non-metal) (prefix)(first syllable of non-metal)-ide.

  • Prefixes indicate the number of atoms of each nonmetal (e.g., mono, di, tri). Covalent naming uses prefixes (mono, di, tri, etc.) and ends with -ide for the second element.

  • Do not add a prefix to the first element if there is only 1 atom.

Atomic Structure and Isotopes

Atomic Mass and Isotopic Composition

• The atomic mass constant is defined as the mass of a carbon-12 atom. Atomic numbers are whole integers, but atomic mass reflects the weighted proportion of different isotopes and is often expressed to decimal places.

• Calculating relative atomic mass from isotopic composition: This involves taking a weighted average of the masses of an element's isotopes, based on their natural abundance.

  • Example: For 99% Carbon-12 (12 amu) and 1% Carbon-13 (13 amu): $(12 \times 0.99) + (13 \times 0.01) = 12.01 \text{ amu}$

  • Example: For Hydrogen isotopes (95% H-1, 4.5% H-2, 0.5% H-3): $(1 \times 0.95) + (2 \times 0.045) + (3 \times 0.005) = 1.055 \text{ amu (approx)}$

• Relative atomic mass is a weighted average of isotopic masses; mass numbers and abundances determine amu values.

Isotopes and Radioisotopes

• Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons.

• Radioisotopes: Isotopes that are radioactive.

• Every element has at least one radioactive isotope; over 3300 radioisotopes are known.

Nuclear Stability

• Radioactivity results from unstable nuclei.

• Radiation is emitted when a radionuclide changes into a more stable nucleus.

• There are 264 stable isotopes among all elements; all isotopes with atomic number (Z) greater than 83 (bismuth) are radioactive.

• Unstable isotopes are influenced by two opposing forces in the nucleus:

  • Electrostatic repulsion between protons.

  • The nuclear force between nucleons (protons and neutrons), which is attractive at short range but repulsive at long range.

• Stability criteria:

  • Small nuclei (Z < 20) are generally stable if the number of protons equals the number of neutrons.

  • Larger nuclei require more neutrons than protons for stability.

  • All nuclei with Z > 83 are unstable.

Radioactivity and Nuclear Decay

Half-life

• Half-life: The time required for the radioactivity level of a sample to be halved, or for half of the mass of a radioisotope to decay.

• It is a characteristic constant for each radioisotope and is used to date materials or determine decay rates.

Types of Radiation

Alpha, beta, and gamma radiation differ in mass, charge, and penetrating power; gamma is highly penetrating; beta is mid-range; alpha is least penetrating but most damaging if ingested.

• Alpha decay: Occurs when a nucleus is too heavy (excess protons and neutrons), emitting an alpha particle (two protons and two neutrons, equivalent to a helium nucleus).

  • Example: Radon decay.

• Beta decay: Occurs when there are too many neutrons for the number of protons. A neutron changes into a proton and an electron, with the emitted electron called a negative beta particle.

  • Example: Cobalt-60 decay.

• Gamma radiation: Highest frequency electromagnetic radiation, emitted by excited nuclei. It often accompanies alpha or beta decay.

Electronic Structure and Light Emission

Atomic Models

• Bohr view: Electrons are tiny, hard particles in specific circular orbits, with electrons in different shells possessing different energies.

• Schrödinger view: Electrons have wave-like properties.

SPDF Notation and Atomic Orbitals

• SPDF notation: Describes how electrons occupy three-dimensional spaces around the nucleus, known as orbitals.

• Energy levels (shells) and subshells: Represented by s, p, d, f.

• Shell capacities:

  • 1st shell: s only

  • 2nd shell: s and p

  • 3rd shell: s, p, d

  • 4th shell: s, p, d, f

• Each subshell holds a specific number of electrons.

• Electron configurations follow principles like the Aufbau principle and Pauli exclusion principle, which explain periodic properties.

Flame Tests and Emission of Light

• Flame test:

  • A metal salt is vaporized in a flame.

  • An outer shell electron absorbs energy and moves to a higher, excited energy level.

  • Being unstable, the electron falls back to a lower energy level, emitting light with characteristic frequencies.

  • Each element produces a unique emission spectrum, correlating to the specific colors observed.

• The color observed in flame tests corresponds to electrons returning to ground state from excited states.

Spectroscopy

• Emission spectrum: A unique pattern of bright lines produced when electrons drop back to their ground energy level.

• Emission spectroscopy: Helps qualitatively identify elements present in a sample (not their concentration) by their characteristic emission lines. Most elements show 4-5 lines, with some being stronger and unique.

• Spectroscopy provides qualitative identification of elements via characteristic emission lines.

Periodicity and Periodic Table Trends

Periodicity

• Periodicity: The regular recurrence of properties as elements are arranged by increasing atomic number.

• This includes trends observed across a period (horizontal row) and down a group (vertical column) in the periodic table.

Periodic Table Structure

• Periods: Horizontal rows, where the number of the period indicates the number of electron shells an element possesses.

• The state of an element at room temperature depends on its melting point (mp) and boiling point (bp):

  • Solid: \text{mp} > 25^\circ\text{C}

  • Liquid: \text{mp} < 25^\circ\text{C} \text{ and } \text{bp} > 25^\circ\text{C}

  • Gas: \text{bp} < 25^\circ\text{C}

• Groups: Vertical columns, elements within a group share similar valence electron counts and properties.

  • Group 1: Alkali Metals (1 electron in outer shell)

    • Soft, can be cut with a knife.

    • Relatively low melting points.

    • Low densities.

  • Group 2: Alkaline Earth Metals (2 electrons in outer shell)

    • Shiny, silvery-white.

    • Somewhat reactive at standard conditions.

  • Group 3: Boron Group (3 electrons in outer shell)

    • Reactivity and metallic characteristics increase down the group.

  • Group 4: Carbon Group (4 electrons in outer shell)

    • Tend to share electrons (form covalent bonds).

  • Group 5: Pnictogens (Nitrogen Group) (5 electrons in outer shell)

    • Reactivity increases down the group.

  • Group 6: Chalcogens (Oxygen Group) (6 electrons in outer shell).

  • Group 7: Halogens (7 electrons in outer shell)

    • Diatomic molecules (e.g., $\text{F}<em>2$, $\text{Cl}</em>2$).

    • Highly reactive non-metals; reactivity increases up the group.

  • Group 8: Noble Gases (Full outer shell)

    • Generally unreactive, but can be made reactive under certain conditions.

• Diatomic elements: F, O, N, halogens, and H.

• Periodic trends reflect systematic changes in atomic structure and electron configuration.

Key Periodic Trends

• Nuclear Charge

  • Nuclear charge: The measure of the attractive force felt between valence electrons and the nucleus. More protons lead to a greater attraction.

• Effective Nuclear Charge

  • Core charge (Effective nuclear charge): The attractive force on valence electrons after accounting for shielding by core electrons. It represents the net positive charge experienced by valence electrons.

  • Calculated as: $\text{Core charge} = \text{Number of protons} - \text{Number of core electrons}$

  • Always positive in neutral atoms.

  • Reasons for trends:

    • Across a period, core electrons remain constant while nuclear charge (number of protons) increases, leading to a stronger pull on valence electrons.

    • Down a group, the number of core electrons increases, leading to greater shielding and a less effective pull from the nucleus on the outermost electrons.

  • Trends:

    • Increases across a period (left to right) due to increasing nuclear charge with constant shielding from core electrons.

    • Remains relatively constant or slightly decreases down a group as additional electron shells increase shielding, largely offsetting the increase in nuclear charge.

  • Effective nuclear charge trends explain how shielding and the number of protons influence the attraction felt by valence electrons.

• Atomic Radius

  • Increases down a group: Due to the addition of more electron shells, valence electrons are further from the nucleus.

  • Decreases across a period (left to right): Due to increasing nuclear charge with a roughly constant number of electron shells, pulling electrons closer to the nucleus.

• Electronegativity

  • Electronegativity: The ability of an atom to attract electrons in a chemical bond.

  • Reasons for trends:

    • Higher core charge and smaller atomic radius generally result in higher electronegativity due to the stronger attraction between the nucleus and bonding electrons.

  • Trend: Decreases down a group and increases across a period.

  • Electronegativity drives bond polarity and reactivity differences among elements.

• Ionisation Energy

  • Ionisation energy: The energy required to remove an electron from a gaseous atom.

  • First ionisation energy: Specifically, the energy to remove one electron from each atom in one mole of gaseous atoms to form one mole of 1+ gaseous ions.

  • Measured in $\text{kJ/mol}$ (Avogadro's number, $N_A = 6.022 \times 10^{23}$).

  • Factors affecting ionisation energy:

    • Distance from the nucleus.

    • Magnitude of nuclear charge.

    • Presence of inner shielding electrons.

    • Completeness of the electron shell.

  • Higher electronegativity generally correlates with higher ionisation energy.

  • Ionisation energy trends reflect increasing attraction in smaller, more charged atoms and shielding effects.

• Reactivity with Water

  • Reactivity depends on an atom's ability to donate, accept, or share electrons.

  • Metals: Tend to donate electrons; lower ionisation energy generally means more reactive.

    • Reaction: $\text{Reactive metal} + \text{water} \rightarrow \text{metal hydroxide} + \text{hydrogen}$

    • Trend: Reactivity generally decreases down a group for metals (less tendency to lose electrons) and increases across a period for metals.

  • Non-metals: Tend to attract electrons; higher electronegativity generally means more reactive with water.

    • Non-metals react by gaining electrons to achieve a stable octet.

    • Trend: Reactivity generally increases towards the right across a period for non-metals.

  • Non-metals typically do not react with water itself, but their reactivity patterns with other elements are influenced by electronegativity.

  • Summary of Reactivity Trends:

    • Group 1 and Group 2 metals: Reactivity with water increases down the group.

    • Transition metals (Groups III and IV): Reactivity generally decreases down their respective groups.

    • Non-metals: Reactivity with water generally decreases down groups due to shielding effects.

    • Across a period: Reactivity is affected by increasing electronegativity and ionisation energy.

    • Noble gases: Generally inert, but their reactivity can increase under certain conditions.

  • Group trends show metals typically become less reactive down the group; non-metals become more reactive across a period.

Chemical Bonding and Molecular Structure

Electronegativity, Valency, and Lewis Diagrams

• Electronegativity and valence electrons are crucial in determining bonding.

• Compounds typically form between atoms with differing electronegativities: if the difference is less than 0.5, the bond is likely covalent.

• The valence shell (outermost shell) contains valence electrons, which are involved in bonding. Atoms gain or lose electrons to achieve a full outer shell.

• Valency: The number of electrons an atom typically uses in bonding.

  • Metals tend to lose electrons, forming cations (positive ions).

  • Non-metals tend to gain electrons, forming anions (negative ions).

  • Noble gases are inert because they already have full outer shells.

• Lewis dot diagrams: Used to represent ions, with the charge shown in the top-right corner. Lewis structures help visualize valence electron transfer and bond formation.

Ionic Compounds

• Nature of Ionic Bonds:

  • Formed when a metallic atom loses electrons to form a cation, and a non-metal atom gains those electrons to form an anion.

  • The electrostatic attraction between these oppositely charged ions creates an ionic bond, leading to the formation of neutral ionic compounds consisting of metal cations and non-metal anions.

• Common Ionic Compounds and Formula Writing:

  • Key polyatomic ions to remember:

    • Ammonium ($\text{NH}_4^+$)

    • Acetate ($\text{C}<em>2\text{H}</em>3\text{O}_2^-$)

    • Hydroxide ($\text{OH}^-$)

    • Nitrate ($\text{NO}_3^-$)

    • Carbonate ($\text{CO}_3^{2-}$)

    • Sulfate ($\text{SO}_4^{2-}$)

    • Phosphate ($\text{PO}_4^{3-}$)

  • Writing chemical formulae:

    • Write the symbol and charge for the two ions that form the compound.

    • "Criss-cross" the charges: write the magnitude of the cation's charge as the subscript for the anion, and the magnitude of the anion's charge as the subscript for the cation.

  • Formula balancing ensures overall charge neutrality; subscripts reflect the stoichiometry needed to balance charges.

• Ionic Networks:

  • Formed by strong electrostatic attraction between ions, creating a 3D lattice structure (e.g., NaCl).

  • Properties: High melting points, hardness, brittleness, and conductivity when molten or in aqueous solutions.

  • Solubility depends on the interactions between ions and water. Ionic networks are brittle.

Covalent Bonding

• Nature of Covalent Bonds:

  • Formed between non-metal atoms through the sharing of electron pairs.

  • Octets are achieved through covalent bonding, which can involve single, double, or triple bonds.

• Polarity of Covalent Molecules:

  • Polar molecules: Have an uneven distribution of electron density, creating dipoles. This occurs when two atoms in a bond have significantly different electronegativities.

  • Nonpolar molecules: Have a symmetric electron distribution where individual bond dipoles cancel out, resulting in zero net dipole.

  • Polar covalent bond: Forms when there's a significant electronegativity difference, but not enough for a complete electron transfer (which would form an ionic bond).

  • Lewis dot diagrams are useful for illustrating polar covalent bonds and identifying lone pairs.

  • The classification of a molecule as polar or nonpolar depends on both electronegativity differences and its molecular geometry; hydrogen bonding is a special case of strong dipole-dipole interaction.

• Molecular Shape (VSEPR Theory):

  • Dipole: An asymmetric charge distribution within a molecule, giving it a distinct positive and negative end. Polar molecules have a net dipole due to this asymmetry, while nonpolar molecules have a zero net dipole.

  • Valence Shell Electron Pair Repulsion (VSEPR) Theory: Predicts molecular shapes by stating that electron pairs around a central atom will arrange themselves to minimize repulsion. Shapes are predicted from Lewis structures.

  • The geometry of a molecule directly determines its polarity. Lone pairs of electrons exert more repulsion and significantly influence bond angles and the overall molecular shape. Use VSEPR to predict geometry from Lewis structures; lone pairs occupy more space and alter angles.

Examples of VSEPR Shapes

• Methane ($\text{CH}_4$): Tetrahedral shape with H–C–H angles of approximately $109.5^\circ$; it has four bonding pairs and no lone pairs on the central carbon.

• Ammonia ($\text{NH}_3$): Pyramidal molecular shape with one lone pair on the nitrogen atom; its electron pair geometry is tetrahedral.

• Water ($\text{H}_2\text{O}$): Bent molecular shape with two lone pairs on the oxygen atom; its electron pair geometry is tetrahedral.

• Carbon dioxide ($\text{CO}_2$): Linear shape with an O=C=O angle of $180^\circ$; for VSEPR, double bonds are treated as single regions of electron density.

• Methanal ($\text{CH}_2\text{O}$): Trigonal planar shape around the central carbon, with three regions of electron density.

Intermolecular Forces

Types of Intermolecular Forces

• Intramolecular forces: Forces within a molecule (covalent, ionic, metallic bonds).

• Intermolecular forces (IMFs): Forces between molecules; much weaker than intramolecular forces. IMFs determine melting and boiling points and are affected by molecular polarity and size.

• Dispersion forces / London forces (weakest): Temporary dipoles in one molecule induce temporary dipoles in neighboring molecules. Present in all covalent molecular compounds; their strength increases with molecular size and polarizability.

• Dipole–dipole forces (middle strength): Attractions between the partial positive ($\delta+$) and partial negative ($\delta-$) ends of polar molecules.

• Hydrogen bonding (strongest IMF): A special, strong type of dipole-dipole interaction that occurs when hydrogen is directly bonded to nitrogen (N), oxygen (O), or fluorine (F). The small size of H and the high electronegativity of N, O, or F increase this attraction.

• Covalent molecular compounds: Characterized by generally low melting and boiling points and typically do not conduct electricity in any state due to the lack of free-moving charged particles.

• The balance of intermolecular forces explains observed melting/boiling points and solubility trends.

Hydrogen Bonding

• Polar molecules generally exhibit higher boiling points than nonpolar molecules of similar molar mass, primarily due to stronger dipole-dipole interactions and the presence of hydrogen bonding (in molecules containing H bonded to N, O, or F).

• Hydrogen bonding is a key factor explaining water's unusually high boiling point and surface tension, and is critical for many biological molecules.

Allotropes and Network Solids

Allotropes

• Allotropes: Different forms of one element in the same physical state that exhibit distinctly different properties due to variations in their atomic arrangement or crystal structure. Allotropy results from different crystal structures; physical properties depend on bonding and arrangement of atoms.

Carbon Allotropes

• Covalent network solids: Atoms linked by a continuous network of covalent bonds throughout the structure.

• Diamond: Each carbon atom is bonded to four others in a 3D tetrahedral lattice; extremely strong, rigid, very hard, and transparent; does not conduct electricity due to localized electrons; very high melting point.

• Graphite: Each carbon atom is bonded to three others in planar sheets, creating a hexagonal lattice. Layers are held together by weak dispersion forces. Each carbon has one delocalised electron, allowing it to conduct electricity within the planes. The weak interlayer forces make the sheets slippery, giving it lubricating properties.

• Fullerenes (e.g., Buckminsterfullerene, $\text{C}_{60}$): Consist of 60 carbon atoms forming a spherical cage. Each carbon is bonded to three others; some delocalised electrons allow conduction. They are nonpolar and can dissolve in nonpolar solvents.

Silicon Dioxide ($\text{SiO}_2$)

• Forms a covalent network lattice where silicon (Si) atoms are bonded to oxygen (O) atoms in a repeating 3D network. The chemical formula $\text{SiO}_2$ reflects this ratio.

• Used by Aboriginal people for spearheads and grinding tools due to its hardness (e.g., Quartz).

Properties of Covalent Networks

• Generally hard, brittle, and have very high melting and boiling points.

• They typically do not conduct electricity because their electrons are localized in bonds (graphite is an exception due to delocalised electrons).

Allotropes of Oxygen

• $\text{O}_2$: The common diatomic form of oxygen.

• $\text{O}_3$ (Ozone): Protects against UV radiation in the upper atmosphere but is toxic at ground level.

Allotropes of Phosphorus

• White Phosphorus ($\text{P}_4$): Highly reactive.

• Red Phosphorus: A more stable chain of $\text{P}_4$ units.

• Black Phosphorus: The most stable form, characterized by buckled layers. Stability and reactivity differ among these forms.

• Allotropy is important for materials science (industrial diamonds, batteries, lubricants, and ozone chemistry).

Metallic Bonding

1. Structure:

   • Consists of a lattice of positive metal ions surrounded by a "sea" of delocalised electrons.

   • Strong electrostatic attraction between the metal cations and the mobile electrons holds the lattice together.

2. Properties of Metals:

   • Solid at room temperature (except mercury).

   • Hard, yet malleable (can be hammered into sheets) and ductile (can be drawn into wires) due to the mobility of delocalised electrons.

   • High melting points.

   • Shiny (lustrous).

   • Excellent conductors of electricity and heat because of the mobile electrons.

• Metallic bonding explains many physical properties of metals, including their ductility and conductivity.