Intro to Atoms and Subatomic Particles

Atoms and Subatomic Structure

• Atoms are the tiny building blocks of matter, a general description that aligns with how scientists view everything around us.
• Historical idea: if you keep dividing matter in half, you would eventually reach something indivisible called an atom.
• Observations in chemistry show that chemical reactions conserve mass regardless of state (solids, liquids, or gases) or reaction type.
• Example discussion of conservation of mass in combustion: burning wood involves fuel and oxygen; products in simple idealized terms are substances like CO and CO₂, with fixed masses for given compositions (e.g., a compound with carbon and oxygen can weigh 28 g or 44 g in the simplified description; there are not intermediate masses in this view). This supports the idea that atoms are conserved during reactions.
• The idea that atoms are made of smaller parts gave rise to the term subatomic.
• Three main subatomic particles (often remembered):
  • Protons
  • Neutrons
  • Electrons

Subatomic Particles: Protons, Neutrons, Electrons

• Proton:
  • Positive charge; represented in diagrams by a proton symbol with a superscript “+” to indicate the positive charge (e.g., p⁺).
  • Relatively massive compared to the electron.
• Neutron:
  • Neutral charge (no net charge, sometimes shown as n⁰ as a superscript to indicate no charge).
  • Located in the nucleus alongside protons; contributes to mass but not charge.
• Electron:
  • Negative charge; located outside the nucleus surrounding it in electron clouds or shells.
  • Much less massive than protons or neutrons; electrons contribute little to the total mass but determine the atom’s volume/size and chemical behavior.
• Nucleus vs. electrons in the atom:
  • Protons and neutrons reside in the nucleus and account for ~99.9% of the atom’s mass.
  • Electrons orbit around the nucleus and are responsible for most of the atom’s volume and chemical properties.
• Size scale and orbit analogy mentioned:
  • If the nucleus were the size of a softball, the first electron orbit would be about half a mile away, illustrating how tiny the nucleus is relative to the overall atomic size.
• Placement of electrons and energy:
  • Electrons are closer to the nucleus when their energy is lower and more strongly attracted to the nucleus; farther away, they are less attracted.
• Practical note on symbols:
  • Protons are shown as positively charged (p⁺); neutrons may be shown with a neutral superscript (n⁰); electrons are denoted by e⁻.
• Modern atom image caveat:
  • The simple picture described here reflects an early 20th-century view (often called the 1913 model) which is useful for intuition but not the full quantum mechanical description.

Scientific Notation and Scale

• Scientific notation is used to handle extremely small or large numbers in atomic-scale measurements.
• Example concept: the proton’s mass is a very small number; you can write numbers like $m_p \approx 1.67 \times 10^{-27} \text{kg}$ to express a tiny mass without writing many zeros.
• The idea discussed: writing numbers in decimal form (e.g., 0.000…087) is impractical; scientific notation compresses zeros and makes calculations easier.
• When dealing with very large numbers (like the velocity of light, or other constants), similar notation avoids long strings of zeros.
• The speaker emphasizes the practicality of using scientific notation for calculations in chemistry and physics.

Protons, Neutrons, Electrons in the Nucleus and Electron Shells

• The nucleus contains protons and neutrons and is the mass center of the atom.
• Electrons are located outside the nucleus and form electron shells or orbitals, which determine chemical behavior.
• The masses: the nucleus accounts for most of the atom’s mass; electrons contribute negligible mass individually but collectively define the atom’s size and its interaction with light and chemistry.
• Charge and mass recap:
  • Proton: charge +1, relatively heavy
  • Neutron: charge 0, similar mass to proton
  • Electron: charge −1, very small mass
• Isotopes and atomic notation are essential for quantifying differences in neutrons while keeping the same protons (same element).

Isotopes and Atomic Notation

• Isotopes differ in the number of neutrons but have the same number of protons (same element).
• Atomic notation basics:
  • Atomic number Z = number of protons (defines the element).
  • Mass number A = Z + N, where N is the number of neutrons.
  • The decimal atomic mass on the periodic table is a weighted average of isotopes’ masses.
• Isotope notation (superscripts and subscripts):
  • General form: $^{A}_{Z}X$ where X is the element symbol, A is the mass number, Z is the atomic number.
  • Example: carbon-14 is denoted as $^{14}_{6} ext{C}$ (A = 14, Z = 6).
• How to find neutrons from notation:
  • $N = A - Z$
• Carbon-14 example from the transcript:
  • For $^{14}_{6} ext{C}$, $N = 14 - 6 = 8$ neutrons.
• The transcript mentions a specific isotope of carbon with eight neutrons (carbon-14) and notes that CO₂ is a molecule with one carbon atom bonded to two oxygen atoms, illustrating isotopes and bonding context.
• Deviations in isotopic composition (e.g., copper with two main isotopes Cu-63 and Cu-65):
  • Copper has two main isotopes: Cu-63 and Cu-65.
  • If the natural abundance is roughly 75% Cu-63 and 25% Cu-65, the weighted average atomic mass would be about 63.5 for copper in that sample.
• The element tin is given as an example with multiple isotopes:
  • Tin has 10 isotopes, illustrating that some elements have many isotopes and weighted averages can be complex.
• Summary points from isotope discussion:
  • The integer on the periodic table represents protons (Z).
  • The decimal atomic mass on the periodic table represents a weighted average of isotopes.
  • The mass number A in notation reflects protons plus neutrons; it is always an integer.
  • The isotope notation $^{A}_{Z}X$ encodes both Z and A.
• Practical take-away: Isotopes change neutron count but not the element’s identity; mass numbers directly relate to neutrons via N = A − Z.
• Climate change context (mentioned): Isotopes will be discussed in relation to climate change later in the course.

Allotropes and Carbon

• Allotropy: the same element can exist in different structural forms with different bonding arrangements.
• Carbon examples:
  • Graphite: layers of carbon atoms arranged in hexagonal sheets; each carbon bonded to three others within a layer; layers held together by weaker interlayer forces; easy to split apart; explains pencil lead where layers slide past each other.
  • Diamond (allotropic form not deeply explained in the transcript, but implied as another allotropic form): a strong three-dimensional network.
• Key point: Changing how atoms are bonded (bonding structure) changes material properties, while keeping the same number of protons and electrons.
• Element naming and symbols:
  • Element symbols come from various languages; e.g., gold is Au from aurum (Latin/Greek roots), which reflects historical discovery and naming conventions.
• The carbon discussion also emphasizes how simple formulas and Lewis structures are used to reason about bonding and molecular structure.

The Periodic Symbol System, Names, and Etymology

• Symbols: one- or two-letter symbols (e.g., Na for sodium) are used in formulas and indicate which elements are present.
• Names vs. formulas: formulas (like NaCl) use symbols and subscripts to indicate the number of atoms; names (sodium chloride) are the common language description of the same substance.
• The two concepts are related but not interchangeable: formulas convey composition and ratios; names convey the substance’s name and often its function or class.
• The discussion notes the difference and promises a small-group activity to practice converting between formulas and names.

Ions, Formulas, and Chemical Nomenclature

• Ion concept: atoms can gain or lose electrons to become charged species.
• Example: Sodium in its neutral form has 11 protons and 11 electrons. If one electron is removed, the ion becomes Na⁺ (positive charge of +1).
• The presentation uses NaCl as an example:
  • Sodium chloride is formed from Na⁺ and Cl⁻ ions combining to make the compound NaCl.
  • The representation NaCl uses symbols from the periodic table with subscripts to indicate the ratio of atoms in the compound.
• Formulas vs. names:
  • The formula NaCl conveys the chemical composition (1 sodium, 1 chlorine).
  • The name sodium chloride conveys the substance’s common or systematic name.
  • There are rules for how formulas map to names and vice versa, which will be practiced in the upcoming small-group activity.
• The transcript foreshadows a small-group activity on Tuesday to practice writing formulas and interpreting names.
• The student is reminded that an element’s symbol (e.g., Na) does not convey information about isotopes or charge on its own; isotopes require separate notation (superscripts and subscripts) and charge requires superscripts (e.g., Na⁺).

Formulas, Names, and Nomenclature in Practice

• The relationship between formulas and names is important for both writing and interpretation:
  • Writing chemical formulas uses the symbols from the periodic table and subscripts to indicate the number of atoms of each element in the molecule or compound (e.g., NaCl).
  • Reading or saying the formula as a name links to the element names and sometimes to the compound’s common name (e.g., sodium chloride).
• The plan to practice this interchangeably indicates the real-world need to be fluent in both representations.

Connections to Broader Themes and Practical Implications

• Foundational principles:
  • Conservation of mass in chemical reactions supports the idea that matter is composed of atoms that are rearranged, not created or destroyed.
  • The existence of subatomic particles and their properties explains chemical behavior, bonding, and material properties.
• Real-world relevance:
  • Isotopes and their weighted abundances are essential for understanding chemistry and climate science; isotopic compositions can serve as tracers in climate studies (to be explored later in the course).
  • Carbon’s allotropic forms illustrate how bonding and structure influence material properties (e.g., graphite in pencils vs. diamond in jewelry or industrial applications).
• Ethical and practical implications:
  • Accurate notations and naming conventions are essential for clear communication in science, safety, and industry.
  • Understanding conservation laws and atomic structure underpins many technologies, environmental studies, and policy decisions.

Quick Reference: Key Equations and Notations from the Transcript

• Atomic number and mass number notation:
  • Z = number of protons (atomic number)
  • A = Z + N (mass number; N = number of neutrons)
  • Isotope notation: $^{A}_{Z}X$ (X is element symbol)
• Neutron count from isotope notation:
  • $N = A - Z$
• Isotope example:
  • Carbon-14: $^{14}_{6} ext{C}$ with $N = 14 - 6 = 8$ neutrons
• Atomic mass concept: decimal atomic mass is a weighted average of isotopes (not a simple integer)
• Ion example:
  • Sodium ion: $ext{Na}^+$ (one electron removed from neutral Na)
• Formula vs. name example:
  • Sodium chloride: formula $ext{NaCl}$; name = sodium chloride
• Simple bonding example note:
  • In carbonate (CO₂), the subscript 2 indicates two oxygen atoms bonded to carbon
• Graphite structure note (conceptual): layers of carbon atoms connected in sheets; each carbon bonded to three others within the layer; interlayer bonds are weak, enabling layering and cleavage

End of Notes