Composition of Matter - Study Notes

Matter: anything that occupies space and has mass.
Mass: the quantity of matter an object has.
Mass vs. weight: weight depends on gravity; mass remains constant; in many contexts we distinguish between the two (W = mg).

Element: pure substance that cannot be broken down into simpler substances by chemical means.
There are more than 100 elements; approximately 30 are important to life.
Six most important elements for life: C, H, O, N, P, S.
Next six important elements: Na, Mg, Cl, K, Ca, Fe.

Atom: the basic unit of a chemical element; the nucleus is the central core.
Subatomic particles:
- Proton: positive charge
- Neutron: neutral
- Electron: negative charge
Nucleus contains protons and neutrons; electrons orbit in energy levels (shells).
Atomic number (Z): number of protons in the nucleus.
Mass number (A): total number of protons and neutrons in the nucleus.
Relationship: $A = Z + N\,.$ (where N is the number of neutrons)
Electron energy levels (shells) and capacities (as shown in the material):
- 1st shell: $2\,e^-\,$
- 2nd shell: $8\,e^-\,$
- 3rd shell: $8\,e^-\,$
Some example electron configurations noted in the table (not exhaustive):
- 2nd shell examples include Li, Be, B, C (with increasing atomic numbers listed in the transcript).
- 3rd shell examples include Na, Mg, Al, Si.

Atomic number: Z = number of protons.
Mass number: A = total number of protons and neutrons in the nucleus.
Example: Phosphorus atom (^{31}_{15}P):
- Protons: 15
- Electrons: 15 (assuming neutral atom)
- Neutrons: 31 − 15 = 16
- Mass number A = 31; Atomic number Z = 15; Neutrons N = 16.

Isotopes of the same element have the same number of protons (chemical behavior is similar), but may have different numbers of neutrons and thus different masses.
Some isotopes are unstable and radioactive; their nuclei decay spontaneously, emitting particles and energy.
Carbon-12 is the most common isotope: (^{12}{6}C), comprising about 99% of naturally occurring carbon; other isotopes include (^{13}{6}C) and (^{14}_{6}C).
Isotopes are important in medicine and research as markers for nonradioactive counterparts.
Radioactive isotopes mentioned: Carbon-14 ((^{14}{6}C)) and Hydrogen-3 (tritium, (^{3}{1}H)).

Compound: substance made of two or more elements.
Chemical reaction: bonds are broken, atoms rearranged, and new bonds are formed.

Covalent bond: sharing of one or more pairs of electrons between atoms.
Molecules can have single, double, or triple covalent bonds.
Representations of molecules include:
- Electron structural diagrams
- Space-filling models
- Structural formulas

(a) Hydrogen (H2): electron distribution and bonding evidence shown via diagram and formulas like H–H.
(b) Oxygen (O2): O=O double bond depiction.
(c) Water (H2O): H–O–H with lone pairs depicted in various models.
(d) Methane (CH4): tetrahedral arrangement with C in the center and four H atoms.

For each molecule, there are several ways to depict structure:
- Electron Structural Diagram
- Structural Formula (showing bonds)
- Space-Filling Model
Examples shown in the material illustrate how H2, O2, H2O, and CH4 are drawn in each representation.

Electronegativity: the tendency of an atom to attract electrons in a bond.
Nonpolar covalent bonds: atoms have equal electronegativity; electrons shared equally.
- Examples: H2, O2, CH4.
Polar covalent bonds: atoms have unequal electronegativity; electrons pulled more toward one atom.
The degree of polarity depends on the difference in electronegativity between the bonded atoms.

Determine nonpolar vs polar covalent bonds:
1) Nitrogen molecule (N2): nonpolar
2) Methane (CH4): nonpolar
3) Ammonia (NH3): polar
4) Formaldehyde (CH2O): polar
Reason: N and O are highly electronegative; C and H have lower and similar electronegativities; greater difference yields polarity.

Ionic bond: transfer of electrons from one atom to another.
Ion: atom or molecule with a net electric charge.
Cation: positively charged ion; Anion: negatively charged ion.
Example: Sodium and Chlorine reaction forming ionic compound:
- Sodium tends to lose an electron: (Na \rightarrow Na^+ + e^-).
- Chlorine tends to gain an electron: (Cl + e^- \rightarrow Cl^-).
- Resulting in sodium chloride: (Na^+ + Cl^- \rightarrow NaCl).
Descriptive illustration in the transcript shows Na and Cl atoms forming NaCl and the ions Na+ and Cl-.

Sodium atom and Chlorine atom shown separately leading to the sodium ion (Na+) and chloride ion (Cl-).
Final salt representation: NaCl.

Role of weak bonds: help determine 3-D shapes and functions of large molecules.
Hydrogen bond: a hydrogen atom covalently bonded to an electronegative atom is also attracted to another electronegative atom.
Van der Waals interactions: weak attractions between nonpolar molecules due to transient positive and negative regions.

Matter, mass, and weight distinctions; mass is inherent to matter, weight depends on gravity.
Elements are fundamental substances; many elements exist, but only a subset is essential for life.
Atoms have a nucleus with protons and neutrons and electrons in energy shells; Z denotes the number of protons, A the total of protons and neutrons, with A = Z + N.
Isotopes differ in neutron number; chemically similar but mass and sometimes stability differ; radioactive isotopes have practical uses in medicine and research.
Molecules form via covalent bonds (sharing electrons) or ionic bonds (transfer of electrons); ions form charged species.
The structure and polarity of bonds determine physical properties and reactivity; electronegativity differences decide bond polarity.
Weak interactions like hydrogen bonds and Van der Waals forces contribute to the shape and behavior of larger molecules.