The History and Evolution of Chemistry and Atomic Theory

Pre-Alchemy and Ancient Theoretical Foundations

Tapputi-Belatekallim

• Role: Overseer (Belatekallim) and Perfume Maker.

• Date: 1200 BC.

• Historical Significance: Documented as an early practitioner of chemical processes through perfume making.

Ancient Greek Philosophers (450 BC)

• The Four Elements: The belief that all matter was composed of four basic "elements":

  1. Earth

  2. Air

  3. Fire

  4. Water

Democritus (460 – 370 BC)

• Basic Concept: Proposed that there are various basic building blocks from which all matter is made.

• Atomic Composition: Everything is composed of small atoms moving with nothing but empty space between them.

• Reception: His ideas were rejected by the leading philosophers of his era.

The Alchemical Era (721 – 1815 AD)

Jabir ibn Hayyan (721–815)

• Transition to Science: Instrumental in moving ideas from the realm of magic to formal science.

• Laboratory Innovation: Created glassware and chemical processes that remain recognized and used in modern chemistry.

Paracelsus (Born 1493)

• Biological Alchemy: Believed he could rearrange the amounts of elements within the human body to change a person.

• Search for Immortality: Desired immortality and searched for a mystical "elixir of life."

Major Contributions of Alchemists

• Discovery of Elements: Identified and isolated the elements mercury, sulfur, and antimony.

• Data Recording: Elements' properties were systematically discovered and recorded.

• Material Science: Created various alloys.

• Laboratory Procedures: Developed fundamental lab procedures, such as the preparation of acids.

Modern Atomic Theory and the Evolution of the Atom

John Dalton (1766–1844)

• Year of Introduction: 1803.

• Key Tenets of Dalton’s Atomic Theory:

  1. Each element is composed of extremely small particles called atoms.

  2. All the atoms of a given element are identical, but they differ from those of any other element.

  3. Atoms may move from one substance to another, but atoms cannot be created, destroyed, or transformed into another element.

  4. Atoms cannot be subdivided.

  5. A given compound always has the same relative numbers and kinds of atoms.

• Dalton’s Model: To account for different properties, he theorized that some atoms are round, pointy, oily, or have hooks.

J.J. Thomson (1856–1940)

• Experiment: Cathode Ray Tube.

• Mechanism: Electricity passes from the cathode to the anode. A small bit of electricity passes through a hole in the anode to strike a phosphor coating.

• Outcomes (1897):

  • Provided proof of subatomic particles.

  • Provided evidence that electrons are negatively charged.

• 1913 Discovery: Identified the existence of isotopes.

• Model: The "Plum Pudding Model."

Ernest Rutherford (1871–1937)

• Experiment (1909): Gold Foil Experiment.

  • Setup: A radioactive source emits a beam of alpha particles through a lead screen with a slit toward a piece of gold foil. A movable fluorescent screen detects where particles land.

  • Observations:

  • Beam A: Transmitted beams with little or no deflection.

  • Beam B: Scattered beams with small deflection.

  • Beam C: Scattered beams with large deflection.

• Outcomes:

  • Discovery of the Nucleus: Characterized as small, dense, and positive.

  • Evidence showed that atoms are mostly empty space.

• Model: The Rutherford Model.

Atomic Data and Isotopes

Element Notation Format

Elements are written in a specific shorthand: ${}^{A}_{Z}X$

• $X$: Element symbol.

• $A$: Mass number (total number of protons and neutrons).

• $Z$: Atomic number (number of protons).

Isotopes

• Definition: Different forms of the same element with different masses.

• Consistency: They have the same number of protons ($p^+$).

• Variance: They have a different number of neutrons ($n^0$).

• Writing Formats:

  • Variant 1: Name-Mass Number (e.g., Carbon-12; Uranium-235).

  • Variant 2: Nuclear Symbol notation (e.g., ${}^{235}_{92}U$ where 235 is the mass number and 92 is the atomic number).

  • Example Structure: Uranium-235 contains $92 \text{ protons}$ and $143 \text{ neutrons}$.

Atomic Mass Standards

• Relative Atomic Mass: The mass of Carbon-12 is defined as exactly $12 \text{ atomic mass units (u)}$. This serves as the defined standard for all other atoms, which are expressed relative to it.

• Average Atomic Mass: A weighted average of the atomic masses of naturally occurring isotopes of an element, as found on the periodic table.

Sample Calculation: Average Mass of Copper ($Cu$)

• Isotope 1: 69.17 \text{%} Copper-63 (Atomic mass: $62.939 \text{ u}$)

• Isotope 2: 30.83 \text{%} Copper-65 (Atomic mass: $64.927 \text{ u}$)

• Calculation: $(0.6917) \times (62.939) + (0.3083) \times (64.927) = 63.546 \text{ u}$

• Extension Example: A ${}^{65}Cu$ isotope with a $+2$ charge contains $29 \text{ protons}$, $36 \text{ neutrons}$, and $27 \text{ electrons}$.

The Periodic Table

Definition and Organization

• Definition: An arrangement of elements in order of their atomic numbers so that elements with similar properties fall in the same column.

• Proton Progression: Each successive element has one more proton than the element preceding it.

• Structure:

  • Groups: Vertical columns.

  • Periods: Horizontal rows.

History: Dmitri Mendeleev

• First Publication: 1869.

• Predictive Power: Left empty spaces in his table for elements not yet discovered (e.g., Scandium). When discovered, they fit his predictions almost perfectly.

Element Families

• Alkali Metals (Group 1): Extremely reactive; not found as free elements in nature (e.g., Potassium).

• Alkaline-Earth Metals (Group 2): Reactive, but less so than Group 1 (e.g., Magnesium, Calcium).

• Transition Elements (Groups 3–12): Metallic (e.g., Copper, Silver, Iron); good conductors; less reactive than Groups 1 and 2.

• Group 13: Boron Family.

• Group 14: Carbon Family.

• Group 15: Nitrogen Family.

• Group 16: Oxygen Family.

• Halogens (Group 17): Most reactive nonmetals (e.g., Bromine, Iodine); react with metals to form salts.

  • Example Reaction: $2K + Cl_2 \rightarrow 2KCl$

• Noble Gases (Group 18): Non-reactive (e.g., Xenon, Krypton).

Classification of Elements

• Metals: Shiny, good electrical conductors.

• Nonmetals: Dull, bad electrical conductors.

• Metalloids: Can be shiny or dull, weak electrical conductors.

Chemical Compounds and Nomenclature

Classification of Compounds

• Molecular (Covalent) Compounds:

  • Created through shared electrons between two atoms (typically two non-metals).

  • Characteristic: Lower melting and boiling points; poor electrical conductors.

• Ionic Compounds:

  • Created through attraction between a cation (+) and an anion (-); electrons are exchanged.

  • Always consist of two ions (monoatomic or polyatomic).

  • Characteristic: High melting and boiling points; good electrical conductivity when molten or in aqueous solution; solids at room temperature; usually form crystals.

Naming Molecular Compounds

1. The element farthest to the left in the periodic table is named first.

2. If both elements are in the same group, the lower one is named first.

3. The suffix -ide is added to the second element.

4. Greek prefixes indicate the number of each element (except "mono-" on the first element).

   • Example: $P_2O_5$ is Diphosphorus pentoxide.

   • Example: $CCl_4$ is Carbon tetrachloride.

   • Example: $N_2O_3$ is Dinitrogen trioxide.

Greek Prefixes List

• 1: mono

• 2: di

• 3: tri

• 4: tetra

• 5: penta

• 6: hexa

• 7: hepta

• 8: octa

• 9: nona

• 10: deca

Diatomic Molecules

Common molecules existing as pairs in their natural state:

• $Cl_2$, $F_2$, $H_2$, $Br_2$, $N_2$, $O_2$, $I_2$

Ion Vocabulary

• Ionic Bond: Strong electrical force between oppositely charged ions.

• Ions: Charged particles resulting from gain or loss of electrons.

  • Anion: Negatively charged.

  • Cation: Positively charged.

• Electrolyte: A solution that can conduct electricity (categorized as Strong, Weak, or Non-Electrolyte).

• Types of Ions:

  • Monatomic: Ion consisting of a single atom (e.g., $Na^+$, $Cl^-$).

  • Polyatomic: Ions with more than one atom (e.g., $SO_4^{2-}$, $NO_3^-$).

Polyatomic Ions to Memorize

• Ammonium

• Acetate

• Dichromate

• Carbonate

• Hydrogen Carbonate

• Perchlorate

• Sulfate

• Nitrate

• Hydroxide

• Phosphate

• Cyanide

Naming Ionic Compounds

• Fixed Charge Metals (Families 1, 2, 13): Metallic ion name + root of nonmetal + -ide.

  • Example: Sodium Chloride ($NaCl$).

  • Example: Aluminum Sulfate ($Al_2(SO_4)_3$).

• Variable Charge Metals (Transition Metals and Family 14): Use Roman numerals to indicate charge.

  • Example: Copper(I) is $Cu^+$.

  • Example: Copper(II) is $Cu^{2+}$.

  • Compound Example: $Cu(NO_3)_2$ is Copper (II) nitrate.

Fundamental Laws of Chemistry

Law of Conservation of Mass

• Principle: Mass is neither created nor destroyed in a chemical reaction.

• Formulaic Representation: $A + B \rightarrow AB$

  • $1 \text{ amu} + 3 \text{ units} \rightarrow 4 \text{ units}$

Law of Multiple Proportions

• Principle: If two or more different compounds are composed of the same two elements, those elements will be in different whole-number ratios.

• Example: Carbon and Oxygen

  • $CO$ (Carbon monoxide): $12 \text{ g C} + 16 \text{ g O}$

  • $CO_2$ (Carbon dioxide): $12 \text{ g C} + 32 \text{ g O}$

  • Oxygen ratio is $16:32$ or $1:2$.

Law of Definite Composition

• Principle: A chemical compound always contains the same elements in the same ratio by mass.

• Example: $H_2O$ always consists of two hydrogen atoms and one oxygen atom ($2 \text{ amu H}$ and $16 \text{ amu O}$). The percent by mass remains constant.

Formula Mass Calculation

To find the formula mass of a compound (e.g., $CaCl_2$):

• $1 \text{ Ca atom} \times 40.078 \text{ u/atom} = 40.078 \text{ u}$

• $2 \text{ Cl atoms} \times 35.453 \text{ u/atom} = 70.906 \text{ u}$

• Total formula mass $CaCl_2$: $110.984 \text{ u}$

Questions & Discussion

DOK1: Name the parts of an atom. Protons, neutrons, and electrons.

DOK2: Atoms are to matter as ______ are to buildings? Bricks (or blocks/foundation units).

DOK1: What is the atomic number? The number of protons in the nucleus of an atom.

DOK1: The identity of an atom is dependent on the ________. Atomic number (or number of protons).

DOK1: What is an isotope? Atoms of the same element that have the same number of protons but different numbers of neutrons.

DOK2: Explain how relative/average atomic mass is different than the mass number? Mass number is the count of protons and neutrons in a single specific atom (a whole number), whereas average atomic mass is a weighted average of all naturally occurring isotopes of that element (usually a decimal).

DOK3: Trace the scientific view of the atom from Democritus to Rutherford explaining how laboratory evidence was key to refining the understanding of this submicroscopic phenomenon.

• Democritus: Conceptualized the atom as a basic building block without experimental proof.

• Dalton: Formulated the first modern atomic theory based on chemical reactions and mass ratios.

• Thomson: Used the Cathode Ray Tube to prove the existence of subatomic particles (electrons), shifting the model from a solid sphere to the Plum Pudding model.

• Rutherford: Used the Gold Foil Experiment to demonstrate that the atom has a central, dense nucleus, refuting the Plum Pudding model and showing atoms are mostly empty space.