Chapter 3 Notes: Properties and Chemical Structure of Matter

Atoms and Molecules

The smallest particle of an element: ATOM.
Atoms of the same or different elements can combine into molecules or formula units through chemical bonding.
Molecules: groups of two or more atoms that are chemically bonded together. Examples include:
- Hydrogen molecule: $ext{H}_2$
- Oxygen molecule: $ext{O}_2$
- Nitrogen molecule: $ext{N}_2$
Water: chemical formula $ext{H}_2 ext{O}$ ; composed of 1 oxygen atom and 2 hydrogen atoms.
Common molecules in the slides: Hydrogen (H₂), Ozone (O₃), Oxygen (O₂), Nitrogen (N₂).
Ions: charged particles that result when neutral atoms lose or gain electrons.
- Chloride ion: $ext{Cl}^-$
- Nitride ion: $ext{N}^{3-}$
- Sodium ion: $ext{Na}^+$
- Vanadium ion: $ext{V}^{4+}$
Example ions: chloride, nitride, sodium, vanadium.
CATION vs ANION
- Cation: positively charged (more protons than electrons).
- Anion: negatively charged (more electrons than protons).
- Charges arise from loss or gain of electrons.
How to form a chemical bond
- Atoms either share or transfer their valence electrons to achieve a more stable configuration.

Valence Electrons and Periodic Context

Valence electrons: electrons located in the outermost energy level (valence shell) of an atom.
Valence shell concepts help explain bonding tendencies.
Example concepts shown in the slides include a Boron atom with valence electrons present in its outer shell.
The periodic table layout shown includes noble gases and groupings such as 1A, 2A, 3A, 4A, 5A, 6A, 7A with labels like H, He, Li, Be, B, N, O, F, Ne, Mg, Al, Si, P, C, Ar; noble gases are highlighted in the context of chemical structure.

Ionic Bonding

There are two main types of bonding discussed:
1) IONIC BONDING – complete transfer of valence electrons from a metal atom to a non-metal atom.
Example of a simple ionic compound: magnesium oxide: $ext{MgO}$ (Mg transfers electrons to O).
Everyday examples of ionic compounds include:
- Sodium chloride: $ext{NaCl}$ (table salt)
- Sodium bicarbonate: $ext{NaHCO}_3$ (baking soda)
- Sodium hydroxide: $ext{NaOH}$ (lye)
- Magnesium sulfate: $ext{MgSO}_4$ (Epsom salt)
- Sodium hypochlorite: $ext{NaOCl}$ (bleach)

Covalent Bonding

Covalent bonding: sharing of electrons between two non-metal atoms.
Key idea: unpaired valence electrons are shared to form a covalent bond, resulting in a covalent molecule.
Example: Water forms when oxygen shares electrons with two hydrogen atoms.
- Water: $ext{H}_2 ext{O}$ ; oxygen and hydrogen form bonds by sharing electrons.
Non-polar covalent bonds: electrons shared equally between two identical atoms (e.g., $ext{N}2, ext{O}2, ext{H}_2$ depending on context; more generally, diatomic homonuclear molecules).
Polar covalent bonds: electrons shared unequally due to electronegativity differences; leads to partial charges on atoms.
Examples of covalent compounds and molecules include water (polar), ozone (O₃), ethanol (C₂H₅OH), carbon dioxide (CO₂), methane (CH₄), and others cited.
Ozone (O₃) is noted as a molecule that protects Earth from UV rays.
Ion formation is not required for covalent bonding; covalent bonding often leads to discrete molecules like H₂O, CO₂, CH₄, etc.
Polar vs non-polar covalent bonds influence molecular properties such as polarity and solubility.

Molecular Shape and Polarity (Lewis Structures and VSEPR)

Lewis structures: dots representing bonding and non-bonding (lone pair) valence electron pairs around atoms.
How to convert a molecular formula to a Lewis structure: example for water, $ext{H}_2 ext{O}$ , which has two H atoms and one O atom.
Lewis structures lead to predictions of molecular shape via VSEPR theory.
VSEPR theory (Valence Shell Electron Pair Repulsion): to obtain molecular shape from the Lewis structure, arrange electron groups around a central atom as far apart as possible to minimize repulsions.
Electron group: a region around an atom that contains electrons; includes lone pairs, single bonds, double bonds, and triple bonds.
Bonding group: single, double, or triple bond.
Non-bonding group: lone pair.
Examples:
- Ammonia: NH₃ — bond pairs = 3, lone pair = 1.
- Water: H₂O — bond pairs = 2, lone pair = 2.
Common molecular shapes (3D arrangement):
- Linear, bent/angular, trigonal planar, tetrahedral, trigonal pyramidal, trigonal bipyramidal, octahedral.
Molecular shape notation: A = central atom; B = surrounding atom or bonding electron group; N = non-bonding valence electron group; x and y = integers referring to the number of bonding or non-bonding electron groups.
- Notation: $ABxNy$

Polarity

Polarity refers to how electrons are shared in a bond or molecule.
Polar bond: electrons are shared unequally (one atom pulls electrons more strongly); results in a slight negative charge on one end and a slight positive charge on the other end (like a tiny magnet).
Non-polar bond: electrons are shared equally between atoms (usually between identical atoms or very similar electronegativities).
Polar molecule: a molecule that has polar bonds or lone pairs that lead to an overall dipole moment.
- Examples of polar bonds include: $ext{H-F}, ext{H-O}, ext{H-N}$ bonds (illustrative).
- Water is a classic polar molecule with the structure $ext{H}_2 ext{O}$ and a bent shape due to lone pairs on oxygen.
Non-polar molecule: a molecule that is nonpolar if it is diatomic with only one kind of atom, or if it consists solely of carbon and hydrogen.
- Examples include diatomic nitrogen $ext{N}2$ , diatomic oxygen $ext{O}2$ , carbon dioxide $ext{CO}2$ , methane $ext{CH}4$ , etc.
Electronegativity plays a key role in determining bond polarity.

Solubility, Solubility in Water, and Miscibility

Solubility: a substance that dissolves and homogeneous mixes with another; dissolution in water depends on polarity.
Solubility in water: polar substances tend to be soluble in water; non-polar substances tend to be insoluble in water.
“Like dissolves like”: polar substances dissolve polar substances; nonpolar substances dissolve nonpolar substances.
Examples of polar substances: water ( $ext{H}2 ext{O}$ ), ammonia (NH₃), ethanol ( $ext{C}2 ext{H}_5 ext{OH}$ ).
Examples of non-polar substances: carbon dioxide ( $ext{CO}2$ ), methane ( $ext{CH}4$ ), benzene ( $ext{C}6 ext{H}6$ ).
Miscibility: the ability of two liquids to mix together.
- Miscible: two liquids that mix in all proportions.
- Immiscible: two liquids that do not mix.
- Relationships by polarity:
- Polar + Polar → Miscible
- Nonpolar + Nonpolar → Miscible
- Polar + Nonpolar → Immiscible
Practical examples:
- Polar pair: water and ethanol are miscible.
- Nonpolar pair: oil and gasoline are miscible.

Summary (Key Concepts)

Atoms are the basic units of elements; when combined, they form molecules or ions.
Valence electrons determine bonding behavior; bonding occurs via sharing (covalent) or transfer (ionic).
Ionic bonding involves complete transfer of electrons from metals to non-metals; examples include NaCl, NaOH, NaHCO₃, MgSO₄, NaOCl, and MgO.
Covalent bonding involves sharing electrons between non-metals; results in discrete molecules such as $ext{H}2 ext{O}, ext{CO}2, ext{CH}4, ext{O}3$ , etc.; bonds can be polar or non-polar.
Lewis structures, along with VSEPR theory, allow prediction of molecular shapes by considering the arrangement of bonding and non-bonding electron groups around a central atom.
Electron group geometry (and the ABxNy notation) helps classify molecular shapes.
Polarity arises from unequal sharing of electrons (polar bonds) or from molecular geometry with lone pairs; non-polar bonds and molecules arise from equal sharing or symmetric structures.
Solubility and miscibility depend on polarity and intermolecular interactions; like dissolves like, and polar vs nonpolar interactions determine whether substances mix.
Real-world relevance includes predicting physical properties (solubility, boiling/melting behavior) and understanding environmental and practical implications (e.g., water’s role as a solvent, ozone’s protection of UV radiation).

Notes on Figures and Notation (from Slides)

Hydrogen, Oxygen, Nitrogen, and Water representations: $ext{H}2, ext{O}2, ext{N}2, ext{H}2 ext{O}$ , $ext{O}_3$ as examples of molecules.
The “ABxNy” notation for molecular shape classification: $ext{AB}{x} ext{N}{y}$ .
Polar vs Non-polar demonstrations include examples like $ext{H-F}, ext{H-O}, ext{H-N}$ bonds (polar) and $ext{N}2, ext{O}2, ext{CO}2, ext{CH}4$ (non-polar) where appropriate.
A note on a displayed formula: a slide shows “F = mc” under Polarity; this appears to be a misprint in the source, as polarity is described via electronegativity and electron sharing, not a force equation. The correct context links polarity to electron distribution rather than a force-mass relationship.