Chapter 1 Notes — Atomic Structure and the Periodic Table (Unit 1)

The Unit Task (Unit 1 Task) focuses on researching a conventional consumer product, its manufacture, and designing a greener version to pitch to investors. Look for Unit Task Bookmarks to connect unit content to the task.
Big ideas (from Unit 1 Overview):
- Every element has predictable chemical and physical properties determined by its structure.
- The type of chemical bond in a compound determines its properties.
- Use of chemicals should be safe to minimize risks to human health and the environment.
Focus on STSE: Green Chemistry explains reducing ecological footprint through safer chemicals, safer processes, and more efficient manufacturing (e.g., safer solvents, fewer steps in synthesis).
Answers to Focus on STSE prompts encourage considering how to green daily life, define green products, and identify industry improvements.

1.1 The Nature of Chemistry

Chemistry: the study of matter and its interactions; matter is anything that has mass and occupies space. Examples include food, cosmetics, batteries, fuels, etc.
Chemistry is often called the central science because it connects physics and biology; many concepts arise from physics (e.g., ionic bonding as attraction of opposite charges).
Chemistry as a human endeavour: scientists design and conduct experiments, collect evidence, and develop explanations/theories; models help visualize the unobservable (microscopic world).
Macroscopic world vs. microscopic world:
- Macroscopic: what we can see, measure, and manipulate directly.
- Microscopic: atoms, ions, molecules; models help visualize these; scientists switch between macroscopic observation and microscopic reasoning.
Chemistry is subdivided into branches: inorganic, organic, nuclear, biochemistry, physical chemistry, etc.
Chemistry vs technology: science explains observations; technology applies those explanations; sometimes science leads to technology (e.g., Teflon) and vice versa; both carry risks and benefits (CFCs and ozone depletion as an historical example).
The international union IUPAC sets standards for chemistry worldwide, promoting international collaboration and ethics in science.
Everyday chemistry: everyone already practices chemistry in daily life (cooking, gardening, hair coloring, etc.).
Theoretical vs empirical knowledge:
- Empirical knowledge comes from observation and experiment.
- Theoretical knowledge arises from models, explanations, and reasoning; both interact to advance chemistry.
Buckminsterfullerene (C60) example: empirical evidence led to a model resembling a soccer ball; fullerene family; illustrates how theory evolves with new data.
Theoretical knowledge evolves over time; heat once thought to be a substance (caloric) replaced by kinetic theory after experiments.
IUPAC and standardization help maintain consistency in naming and chemical practices.
Ending takeaway for 1.1: Chemistry is about matter and interactions; it blends macroscopic experiments with microscopic models; the periodic table and theories arise from and guide experimentation.

1.2 Atomic Structure

The atom as the fundamental unit of matter:
- The oldest idea of the atom dates to Democritus (atomos: indivisible).
- Aristotle argued matter consists of earth, air, water, fire; this held sway for centuries.
- Alchemists contributed lab techniques and apparatus that later informed modern chemistry (glassware, alloys, safety practices).
- Dalton (1808) proposed the Daltonian solid-sphere model: atoms are indivisible, atoms of an element are identical, atoms combine to form compounds, and atoms rearrange in reactions but are not created/destroyed.
Key experiments that shaped atomic theory:
- Thomson’s cathode-ray tube experiments revealed the electron; proposed the plum pudding model (electrons embedded in a positively charged sphere).
- Rutherford’s gold foil experiment: most alpha particles passed through, but some were deflected; concluded a tiny, dense, positively charged nucleus exists with electrons surrounding it in mostly empty space; this led to the nuclear model.
- Chadwick’s discovery of neutrons in the nucleus (1932).
- Bohr’s energy-level model for hydrogen: electrons occupy definite energy levels; emission lines depend on electron transitions between levels; introduced the planetary model with quantized orbits.
Bohr–Rutherford synthesis: a simplified atom model combining Dalton/Thomson/Rutherford/Bohr ideas to visualize nucleus (protons + neutrons) and electrons in energy levels.
Subatomic particles and their characteristics:
- Electron: symbol e−; location: outside the nucleus; charge: −1; mass ~ 9.11×10^-31 kg.
- Proton: symbol p+; location: in the nucleus; charge: +1; mass ~ 1.67×10^-27 kg.
- Neutron: symbol n0; location: in the nucleus; charge: 0; mass ~ 1.67×10^-27 kg.
Notation and notation concepts:
- Atomic number Z = number of protons in the nucleus.
- Mass number A = total number of protons and neutrons (A = Z + N).
- For a neutral atom, Z equals the number of electrons.
- Atomic mass unit (u): defined as 1/12 the mass of a carbon-12 atom;
  $1 \,\mathrm{u} = \frac{1}{12} m\left(^{12}\mathrm{C}\right)$
  equals $1\,\mathrm{u} = 1.660\,540\times 10^{-27}\,\mathrm{kg}$
- Isotopes are atoms of the same element with different numbers of neutrons and thus different mass numbers A.
Notation for atoms: $^{A}{Z}\mathrm{X}$ (e.g., $^{12}{6}\mathrm{C}$ ). The notation gives A (mass number) and Z (atomic number).
Bohr–Rutherford diagrams: use concentric energy level shells; place protons and neutrons in the nucleus; electrons in shells; relevant for the first 20 elements.
The periodic table: organized by increasing atomic number; groups (vertical) share similar properties; periods (horizontal) reflect repeating patterns; periodic law: properties recur periodically with increasing atomic number.
Summary of 1.2: Atomic theory evolved through models; Bohr–Rutherford model remains useful for predicting trends; atomic number Z defines identity; mass number A and neutron count N derived from A − Z; atoms can be represented by multiple notations; atomic mass unit frames measurements for tiny masses.
1.2 Key Equations and Notation
- Mass number: $A = Z + N$
- Number of neutrons: $N = A - Z$
- Atomic notation: $^{A}_{Z}X$ where Z = protons, A = protons + neutrons.
- Atomic mass unit (definition): $1\,u = \tfrac{1}{12} m(^{12}\mathrm{C})$ and the kg equivalent $1\,u = 1.660\,540\times 10^{-27}\ \mathrm{kg}.$
- Bohr–Rutherford diagram conventions: shells represent energy levels; nucleus contains p+ and n0; electrons occupy valence shells; valence electrons defined as electrons in the outermost shell.
1.2 Summary points
- Atomic theory has evolved from indivisible spheres to complex nucleus + electron structure models; the Bohr–Rutherford model remains a practical visualization for many elements.
- Atomic number Z uniquely identifies an element; mass number A identifies a specific isotope; neutrons N = A − Z.
- Atomic mass unit is a relative scale based on carbon-12; mass of atoms is usually expressed in amu rather than kilograms.
- Three subatomic particles: electron (e−), proton (p+), neutron (n0); charges and masses noted above.

1.3 Ions and the Octet Rule

Noble gases are exceptionally stable (inert) and typically do not form compounds.
Octet rule: atoms tend to attain a full valence shell of 8 electrons (except H and He, which are stable with 2 in the valence shell).
How atoms achieve octets (three pathways):
- Covalent bonding (sharing electrons) to complete an octet.
- Ionic bonding by losing electrons to form cations (positive ions) or gaining electrons to form anions (negative ions).
- In some cases, atoms achieve octets by gaining or losing electrons to form ions; the resulting ion has a charge (valence).
Ions and charges:
- Cation: positively charged ion formed when an atom loses electrons (common with metals; located left of the zigzag line in the periodic table).
- Anion: negatively charged ion formed when an atom gains electrons (common with nonmetals; right of the zigzag line).
- Example ions:
- Na → Na⁺ (loss of 1 electron; valence +1)
- Al → Al³⁺ (loss of 3 electrons; valence +3)
- Cl → Cl⁻ (gain of 1 electron; valence −1)
- S → S²⁻ (gain of 2 electrons; valence −2)
Polyatomic ions: ions composed of more than one atom (e.g., NO₃⁻, PO₄³⁻); they behave as a single ion.
Multivalent elements: some metals can form more than one stable ion (e.g., Cu⁺, Cu²⁺; Fe²⁺, Fe³⁺).
- Traditional names: cuprous (Cu⁺), cupric (Cu²⁺); ferrous (Fe²⁺), ferric (Fe³⁺).
- IUPAC nomenclature uses Roman numerals to indicate the oxidation state: copper(I) vs copper(II).
Memory aids:
- Cations: positive, “t” resembles a plus sign; anions: “are negative ions” (A for anion).
- The word “anion” resembles “on-ion,” highlighting its negative charge.
Safety and lab practice notes accompany ion-related investigations in context of Chapter 1 labs.
1.3 Key Equations and Concepts
- Octet rule concept: aim for 8 valence electrons (2 for H/He).
- Ion formation: metals tend to lose electrons to form cations; nonmetals tend to gain electrons to form anions.
- Ionic charges reflect electron transfer; e.g., Na⁺, Cl⁻, Mg²⁺, Al³⁺.
- Polyatomic ions are treated as single ions (e.g., NO₃⁻, PO₄³⁻).
- Multivalent metals have multiple common oxidation states; naming uses Roman numerals in IUPAC conventions.
1.4 Isotopes, Radioisotopes, and Atomic Mass
Isotopes: same Z (protons) but different N (neutrons) and A (mass number).
Isotopic abundance: the relative amounts of isotopes in a sample, usually given as percentages.
Mass spectrometry: a key tool to identify isotopes and their abundances by separating ions by mass-to-charge ratio.
Radioisotopes: isotopes that are radioactive and decay, emitting nuclear radiation (alpha, beta, gamma).
Nuclear decay types:
- Alpha (α): helium-4 nucleus (2p+, 2n0), charge +2; blocked by paper.
- Beta (β): high-energy electron; can pass through paper but is blocked by aluminum.
- Gamma (γ): high-energy electromagnetic radiation; very penetrating; blocked by lead.
Medical uses: radioisotopes like I-131 (thyroid), Tc-99m (imaging).
Isotopes and atomic mass:
- Atomic mass is the weighted average of the masses of naturally occurring isotopes, weighted by isotopic abundance.
- Notation of atomic mass: typically given in amu (atomic mass units).
Isotopic abundance example: natural Mg contains Mg-24 (78.7%), Mg-25 (10.1%), Mg-26 (11.2%).
- Atomic mass calculation: $\text{Atomic mass} = 0.787\times 24 + 0.101\times 25 + 0.112\times 26 = 24.3\,\mathrm{u}$
Mass spectrometer components: ion source, analyzer, detector; used to determine masses and abundances.
Practical application: atomic mass of elements in the periodic table is a weighted average.
1.4 Key Equations
- Isotope notation: $^{A}_{Z}\mathrm{X}$ with Z protons, A mass number, X element symbol.
- Atomic mass calculation (weighted average):
 $\text{Atomic mass} = \sumi\left( fi \cdot Mi \right) = \sumi\left( \frac{\%i}{100} \cdot Mi \right)$
- For a sample with 3 isotopes: Mg example above.
- Mass of isotopes and their abundances summarized in tables; isotopes can be stable or radioisotopes.
- Nuclear decay types and their effects on A and Z (brief): Alpha decay decreases A by 4 and Z by 2; Beta decay increases Z by 1 (A unchanged) for neutron-to-proton conversion; Gamma decay leaves A and Z unchanged.
1.5 The Periodic Table and Periodic Law
The periodic table organizes all known elements by increasing atomic number Z; elements with similar properties align in columns (groups) and rows (periods) reveal periodic trends.
Groups (1–18) and periods (1–7); noble gases (Group 18) are inert; halogens (Group 17) are highly reactive nonmetals; alkali metals (Group 1) are highly reactive metals; alkaline earth metals (Group 2) are reactive metals; transition metals are in the middle.
The concept of representative elements includes Groups 1, 2, and 13–18; transition metals lie in the center.
Lewis symbols and Bohr–Rutherford diagrams are two ways to represent atoms and their valence electrons.
Periodic law: when elements are arranged in order of increasing atomic number, their properties recur in a periodic pattern.
History of the periodic table:
- Newlands’ Law of Octaves (arranged in groups of seven; limited applicability).
- Mendeleev (1869) created a periodic table with gaps, predicting undiscovered elements and their properties.
- Meyer independently developed a similar table; Moseley later refined order by atomic number, solving inconsistencies.
- Seaborg added actinide and lanthanide series and contributed to the understanding of transuranic elements.
IUPAC now standardizes element naming; temporary names exist while discoveries are confirmed (e.g., ununtrium for element 113).
1.5 Key Equations/Notations
- Periodic law: properties recur in a systematic way when elements are ordered by atomic number Z.
- Notation and diagram tools: Bohr–Rutherford diagrams show electron shells; Lewis symbols show valence electrons as dots around the symbol.
1.5 Summary highlights
- The periodic table is a dynamic tool shaped by experimental discoveries and theoretical advances; Lewis symbols complement Bohr–Rutherford diagrams for valence electrons.
1.6 Chemistry Journal: A not-So-Elementary Task
- A segment focusing on the historical and conceptual evolution of chemistry through case studies and activities. (Content described in the unit but not required for core definitions.)
1.7 Periodic Trends in Atomic Properties
Atomic radius (size of atoms) is defined as the distance from the nucleus to the outermost electrons; for diatomic molecules, radius is half the distance between nuclei.
Trends:
- Atomic radius increases down a group (more energy levels added; outer electrons are farther from nucleus).
- Atomic radius decreases across a period from left to right (increasing nuclear charge Z with similar shielding; electrons are pulled closer to the nucleus).
- Noble gases generally have the smallest radii in their periods due to full valence shells.
Ionic radius concepts:
- Cations (positive ions) are smaller than the corresponding neutral atoms because electrons are removed, increasing effective nuclear charge on remaining electrons.
- Anions (negative ions) are larger than their neutral atoms due to added electrons increasing electron-electron repulsion while nuclear charge remains fixed.
Ionization energy (IE): energy required to remove an electron from a gaseous atom or ion.
- First IE: energy to remove the outermost electron.
- Trends:
- IE tends to increase across a period (due to increasing nuclear charge and decreasing radius; tighter hold on electrons).
- IE tends to decrease down a group (radius increases; outer electrons are farther from the nucleus and more weakly held).
- Examples from the text:
- Lithium IE1 ≈ 520 kJ/mol; Calcium IE1 ≈ 590 kJ/mol; Bromine IE1 ≈ 1140 kJ/mol; Chlorine EA ≈ 349 kJ/mol (note: EA is energy released when adding an electron).
Electron affinity (EA): energy change when an electron is added to a neutral gaseous atom.
- Generally more negative across a period (greater tendency to gain electrons) and becomes less negative down a group.
- Helium, Be, N, Mg, Ar show near-zero or positive EA values, meaning energy is required or not released for adding an electron.
Theoretical explanations: effective nuclear charge and radius influence IE and EA; as Z increases across a period with little additional shielding, electrons are held more tightly; as you go down a group, inner shells shield outer electrons, reducing effective pull.
1.7 Periodic Trends in Atomic Properties: summary points
- Atomic radius, ionic radius, ionization energy, and electron affinity exhibit periodic trends explained by effective nuclear charge and shielding.
- Across a period: radius decreases; IE and EA generally increase.
- Down a group: radius increases; IE and EA generally decrease.
- Understand how radius changes relate to IE/EA trends.

1.7 Practice and Investigations (Concepts and Applications)

Investigation 1.4.1: The Nuts and Bolts of Atomic Mass – a hands-on model using nuts and bolts to represent protons and neutrons and to derive atomic mass from isotopic abundances.
Investigation 1.5.1: The Search for Patterns – examining reactivity of metals in water and acids to identify trends by group and period.
Investigation 1.7.1: Graphing Periodic Trends – building graphs of atomic radius and first IE against atomic number to observe trends.
Correlational graphing (Investigation 1.7.1) – comparing atomic radius and first IE across the first 20 elements; analyze two trends on a single graph with appropriate labeling and interpretation.

1.7 Key Formulas and Equations (in-context)

Atomic radius and ionic radius concepts do not have a single simple formula, but the trends are explained by effective nuclear charge (Z_eff):
- As you move across a period, Z_eff increases; outer electrons are pulled closer, reducing radius.
- As you move down a group, additional energy levels increase the average distance of outer electrons from the nucleus, increasing radius.
Ionization energy (IE) and electron affinity (EA) relationships with atomic radius and Z_eff:
- Across a period: IE generally increases; EA increases for nonmetals (more negative EA across a period).
- Down a group: IE generally decreases; EA becomes less negative down a group.

1.7 Quick Reference: Notation and Concepts

Atomic number: Z = number of protons; represented in the periodic table cell; for neutral atoms, Z equals electron count.
Mass number: A = Z + N; Neutrons: N = A − Z.
Isotopes: same Z, different N; isotopic abundance used to calculate weighted atomic mass.
Atomic mass unit: defined relative to C-12; mass of a single atom is expressed in amu; 1 u = 1.660540×10^-27 kg.
Notation for isotopes: $^{A}_{Z}\mathrm{X}$ .
Periodic law: the properties of elements show periodic recurrence when arranged by atomic number; trends across periods and down groups.

Key Terms (condensed glossary)

matter, empirical knowledge, theoretical knowledge, theory, atom, electron, proton, neutron, energy level, valence shell, valence electron, atomic number Z, mass number A, atomic mass unit u, octet rule, ion, cation, anion, multivalent, polyatomic ion, isotope, isotopic abundance, mass spectrometer, radioactive decay, nuclear radiation, alpha particle, beta particle, gamma ray, radioisotope, atomic mass, metalloid, group, period, periodic law, Lewis symbol, atomic radius, effective nuclear charge, ionic radius, ionization energy, electron affinity.

Equations and LaTeX references

Atomic mass unit definitions:
$1\,\mathrm{u} = \frac{1}{12} m\left(^{12}\mathrm{C}\right) \quad\text{and}\quad 1\,\mathrm{u} = 1.660\,540\times 10^{-27}\,\mathrm{kg}.$
Mass number and neutron calculation:
$A = Z + N, \quad N = A - Z.$
Isotopic notation: $^{A}{Z}\mathrm{X}$ (example: $^{12}{6}\mathrm{C}$ ).
Weighted atomic mass calculation (isotopic abundances):
$\text{Atomic mass} = \sumi \left( \frac{\%i}{100} \cdot M_i \right)$
Example (Mg): $0.787\times 24 + 0.101\times 25 + 0.112\times 26 = 24.3\,\mathrm{u}.$
Electron/ion notation:
- Ion charge examples: Na⁺, Cl⁻, Mg²⁺, Al³⁺.
Ionization energy trend (qualitative): IE increases across a period, decreases down a group; numerical examples include IE1(Li) ≈ 520 kJ/mol, IE1(Ca) ≈ 590 kJ/mol, IE1(Br) ≈ 1140 kJ/mol, EA(Cl) ≈ 349 kJ/mol.
Electron affinity examples: Cl has EA ≈ 349 kJ/mol; noble gases have near-zero or positive EA values.
Plan for green chemistry: minimize waste, reduce energy use, select safer solvents; examples include replacing toxic solvents with safer alternatives like liquid CO2.

Connections and implications

The atom’s structure explains elemental properties and periodic trends; electrons in valence shells govern bonding type and reactivity.
The octet rule guides predictions of ion formation and bond formation; valence electron counts explain group properties (alkali metals with 1, alkaline earth with 2, halogens with 7, noble gases with 8).
The periodic table is more than a catalog: it encodes predictive relationships (group/period trends) that guide chemistry in real life (e.g., predicting reactivity, bonding, and compound formation).
Ethical and practical implications: safety in chemical use, environmental considerations (green chemistry), responsible handling of radioactive materials, and the societal relevance of the scientific enterprise (IUPAC standards, disclosure of safety concerns, etc.).

Quick study prompts (to check understanding)

Why do atoms form ions, and how does the octet rule relate to ion formation? Give examples for Na, Cl, and S.
How is the atomic mass of an element determined from its isotopes? Use Mg as an example with given abundances.
Explain the difference between a cation and an anion with Bohr–Rutherford diagrams.
What is the periodic law, and how do periods and groups relate to element properties?
Describe how effective nuclear charge explains atomic radius trends across a period and down a group.
Give an example of a green chemistry approach and discuss how it reduces environmental impact.

// End of Chapter 1 Notes