CHEM051 - CH4 PT 1

Elements, Atoms & Chemical Symbols

• “Element” ≈ “atom” at its most reduced, indivisible (chemical) form.
• Every element/atom contains a dense nucleus (protons + neutrons) with electrons moving in surrounding space.
• Number of protons (the atomic number, $Z$) uniquely defines the element.

Origin of Element Names

• Named after:
– Planets (Mercury)
– Mythological figures (Titanium ⇾ Titans)
– Minerals/Colors (Chlorine ⇾ Greek “chloros,” green-yellow)
– People (Curium ⇾ Marie Curie)
– Places (Californium ⇾ California)

Chemical Symbols (IUPAC rules)

• 1 or 2 letters only.
– First letter always capital.
– Second (if present) lowercase.
• Purpose = prevent ambiguity (e.g. Co ≠ CO).
• Latin/Greek roots preserved when modern letter conflicts existed:
– Sodium ⇾ “natrium” → Na.
– Potassium ⇾ “kalium” → K.
– Iron ⇾ “ferrum” → Fe.

Diatomic Elements (“Rule of Seven”)

• Seven elements bind as $X<em>2$ in elemental form (for stability): H$2$, N$2$, O$2$, F$2$, Cl$2$, Br$2$, I$2$.
• Mnemonic: start at N (atomic #7) on the periodic table, trace an “7-shaped” path across to F then down to I; then add H.

Periodic Table Framework

• Columns = Groups/Families; rows = Periods (energy levels).
• Main-group numbering used in this course: $1A, 2A, 3A\dots 8A$.
• Special family names:
– Alkali metals (1A) – extremely reactive, 1 valence e⁻.
– Alkaline-earth metals (2A) – reactive, 2 valence e⁻.
– Halogens (7A) – highly reactive non-metals.
– Noble gases (8A) – inert/stable, complete outer shell.
• Regions:
– Blue (left) = Metals.
– Green stair-step = Metalloids (B, Si, Ge, As, Sb, Te, At).
– Yellow (right) = Non-metals.

Metal vs Non-metal Properties

Historical Development of Atomic Theory

• Democritus (~400 BC) – coined “atomos,” indivisible units; purely philosophical.

Dalton’s Soccer-Ball Model (1808)

1. Elements are composed of tiny, indestructible atoms.

2. Atoms of one element differ from those of another.

3. Atoms combine in simple, whole-number ratios to form compounds (e.g. $\text{H}<em>2\text{O}$ not $\text{H}</em>{2.5}\text{O}$).
   • Model: uniform, featureless solid sphere.

Thomson’s Plum-Pudding Model (1897-1903)

• Cathode-ray-tube experiment → discovered electron (negatively charged).
• Proposed sphere of positive “pudding” with electrons as embedded “plums.”

Rutherford’s Nuclear Model (1910)

• Gold-foil experiment (α-particle scattering).
– Most α passed → atom mostly empty space.
– Few deflected/back-scattered → small, dense, positively charged nucleus.
• Electrons orbit the nucleus; model introduces protons.

Chadwick’s Neutron (1932)

• Discovered neutral nuclear particle → neutron; explains additional mass without charge.

Modern Atomic Structure Summary

• Subatomic particles:
– Proton: charge $+1$, mass ≈ 1 amu.
– Neutron: charge $0$, mass ≈ 1 amu.
– Electron: charge $-1$, mass ≈ ${1\over1836}$ amu.
• Typical radius ratios: atom ≈ $10^{-10}\text{ m}$, nucleus ≈ $10^{-15}\text{ m}$.
• Overall neutral atom: #e^- = #p^+.

Atomic Number, Mass Number & Isotopes

• Atomic number $Z$ = protons.
• Mass number $A$ = protons + neutrons.
• Neutrons: $n = A - Z$.
• Isotopes = same $Z$, different $n$ (thus different $A$).
– (^1\text{H}), (^2\text{H}) “deuterium,” (^3\text{H}) “tritium” are isotopes of hydrogen.
• Ions: unequal p⁺ & e⁻.
– Loss e⁻ → cation (positive).
– Gain e⁻ → anion (negative).

Average (Weighted) Atomic Mass

• Definition: weighted mean of naturally occurring isotopes.
• Formula: $\bar{A} = \sum<em>{i} A</em>i\times f<em>i$ where $f</em>i$ = fractional abundance.
• Example (lecture extrapolation):
Nitrogen isotopes $^{14}\text{N}\, (50\%)$, $^{15}\text{N}\,(50\%)$ → $\bar{A}=14\times0.5+15\times0.5=14.5\,\text{amu}$.
• Closest isotope by % dominantly influences listed periodic-table mass.

Representative Worked Problem (from lecture)

1. Write full isotope notation for chlorine-35: $^{35}_{17}\text{Cl}$ (Z = 17, A = 35).

2. Find neutrons: $n = 35-17 = 18$.

3. Identify isotopes in a set: same $Z$ but different $A$ (exercise with unknowns R & Y yielding identical $Z$).

4. Determine most abundant boron isotope using periodic mass ≈ 10.8 amu → closer to 11 (so $^{11}\text{B}$ dominates).

Connections & Relevance

• Reactivity patterns (e.g., alkali metals in water) derive from easy electron loss predicted by group position.
• Rule of Seven gases dominate Earth’s atmosphere/industry (O$2$, N$2$, Cl$_2$ production, etc.).
• Semiconductor technology exploits metalloid duality (Si, Ge).
• Understanding isotopes underpins radiometric dating, nuclear medicine (e.g., (^3\text{H}) tracing).
• Philosophical arc: shift from speculative atomos (Democritus) to empirical, instrument-based science, illustrating the scientific method’s ethical imperative of evidence over conjecture.

Key Equations & Symbols (Quick Reference)

• Atomic number: Z = #\,\text{protons}.
• Mass number: $A = p + n$.
• Neutrons: $n = A - Z$.
• Average atomic mass: $\bar{A}=\sum A<em>i f</em>i$.
• Charge of ion: $q = p - e$ (units of elementary charge).

Study Tips

• Memorize the 7 diatomic elements; draw the “7” on a blank periodic table.
• Know four special group names & region colors (metal, non-metal, metalloid).
• Associate models with scientists:
Democritus → atomos; Dalton → soccer ball; Thomson → plum pudding; Rutherford → nuclear; Chadwick → neutron.
• Practice converting isotope symbols, calculating neutrons, and weighted masses.

The 7 diatomic elements that bind as $X<em>2$ in their elemental form for stability are: Hydrogen ($H</em>2$), Nitrogen ($N<em>2$), Oxygen ($O</em>2$), Fluorine ($F<em>2$), Chlorine ($Cl</em>2$), Bromine ($Br<em>2$), and Iodine ($I</em>2$). A mnemonic to remember them is to start at Nitrogen (atomic #7) on the periodic table, trace a "7-shaped" path across to Fluorine then down to Iodine, and then add Hydrogen.

The 4 special group names are:

• Alkali metals (1A)

• Alkaline-earth metals (2A)

• Halogens (7A)

• Noble gases (8A)

The region colors are:

• Blue (left) = Metals.

• Green stair-step = Metalloids (B, Si, Ge, As, Sb, Te, At).

• Yellow (right) = Non-metals.

Converting isotope symbols for each element of the periodic table involves using the notation $^{A}<em>{Z}\text{X}$, where $A$ is the mass number (protons + neutrons), $Z$ is the atomic number (number of protons), and $\text{X}$ is the chemical symbol for the element. While it's not feasible to list every isotope for every element, the process remains consistent. For example, chlorine-35 is written as $^{35}</em>{17}\text{Cl}$, where 17 is the atomic number (Z), and 35 is the mass number (A). To find the number of neutrons, you subtract the atomic number from the mass number ($n = A - Z$).

Here's a summary of what was explicitly mentioned to be memorized or practiced:

• Memorize the 7 diatomic elements; draw the “7” on a blank periodic table.

• Know four special group names & region colors (metal, non-metal, metalloid).

• Associate models with scientists:

  • Democritus → atomos

  • Dalton → soccer ball

  • Thomson → plum pudding

  • Rutherford → nuclear

  • Chadwick → neutron

• Practice converting isotope symbols, calculating neutrons, and weighted masses.