Chapter 2: Atoms and Elements
Chapter 2: Atoms and Elements
2.3 Modern Atomic Theory
Core Understanding: The theory that all matter is composed of atoms emerged from observations and scientific laws.
Laws supporting atomic theory:
Law of Conservation of Mass: In a chemical reaction, matter is neither created nor destroyed; the total mass of substances remains constant.
Law of Definite Proportions: All samples of a given compound have consistent proportions of constituent elements, regardless of source or preparation method.
Brownian Motion: Observes the random movement of small particles suspended in fluids, indicating collisions with other fast-moving particles.
2.4 Hydrogen Sulfide Formation Example
Chemical Reaction Example:
Hydrogen and sulfur react to form hydrogen sulfide
Reaction details:
4.03 g of hydrogen reacts with 62.13 g of sulfur.
Total mass of products = mass of reactants = 4.03 g + 62.13 g = 66.16 g of hydrogen sulfide formed.
John Dalton and the Atomic Theory
Key Points of Dalton’s Atomic Theory:
Each element comprises tiny, indestructible particles called atoms.
All atoms of a given element share the same mass and properties, distinguishing them from other elements.
Atoms combine in simple, whole-number ratios to create compounds.
Atoms cannot change into different types of atoms through chemical reactions; they only rearrange.
2.5 The Discovery of the Electron
J. J. Thomson's Research:
Conducted cathode ray experiments that revealed particles called cathode rays originated from the cathode to the anode.
Discovered that these particles, electrons, carry a negative charge.
Thomson's findings proved that electrons are independent of the cathode material used.
Robert Millikan's Oil Drop Experiment:
Millikan determined the charge of a single electron to be coulombs.
2.6 The Structure of the Atom
Thomson's Plum-Pudding Model: Proposed that electrons are embedded in a positively charged sphere.
Rutherford's Gold Foil Experiment:
Investigated the atomic structure by directing alpha particles at a thin gold foil.
Findings: Most particles passed through, but some were deflected or bounced back, revealing that atoms consist of dense cores (nuclei) surrounded by empty space.
Conclusion: Atoms have large regions of empty space with small, dense nuclei.
2.7 Subatomic Particles
Components of Atoms:
Protons: Positively charged particles in the nucleus.
Neutrons: Neutral particles also in the nucleus.
Electrons: Negatively charged particles that orbit the nucleus.
Charges: Protons' (+1.6 x 10^-19 coulombs) and electrons' (-1.6 x 10^-19 coulombs) charges are equal in magnitude but opposite in sign.
Atomic Identity and Numbering
The atomic number (Z) indicates the number of protons in the nucleus, uniquely defining each element.
The mass number (A) is the total number of protons and neutrons in the nucleus (A = Z + N).
2.8 Isotopes
Defined as variants of an element with differing numbers of neutrons.
Example: Two boron atoms: both have 5 protons but different neutron counts.
Isotope Notation:
Mass number (A) = number of protons (Z) + number of neutrons (N).
For boron-10: Mass number = 10, Atomic Number = 5 (Neutrons = 10 - 5 = 5).
2.9 Ions and Atomic Charge
Atoms are electrically neutral with equal numbers of protons and electrons.
Ionic Formation:
Ions form when atoms gain or lose electrons.
Anions: Formed by gaining electrons (negative charge).
Cations: Formed by losing electrons (positive charge).
2.10 The Periodic Table and Periodic Law
Structure: Organized into columns (groups) and rows (periods). Groups exhibit similar chemical properties.
Classification of Elements:
Metals: Good conductors, malleable, ductile, shiny. Located in the lower-left and middle of the table.
Nonmetals: Varied properties, located in the upper right (e.g., C, P, S, Se, I).
Metalloids: Exhibit mixed properties and are located along the zigzag line separating metals and nonmetals (e.g., Si, Ge).
2.11 Usage of Molar Mass
Definition: Molar mass expresses the mass of one mole of a substance, numerically equivalent to the atomic mass of the element in grams per mole.
Conversion:
Mass (in grams) and number of moles are interconversable using molar mass as a conversion factor.
Example: 1 mole of iron (Fe) equals the atomic mass of iron in grams.
Applications: Counting the number of atoms in a sample, determining mass, and calculating atomic quantities based on experimental data.
2.12 Problem Solving: Calculating Atomic Mass and Moles
Example Calculation: To find the average atomic mass of chlorine considering its isotopes and natural abundances:
Mole Representations: Using Avogadro's number, which is , to express large quantities.
Numerical comparisons illustrating the vastness of a mole (e.g., years spent eating hotdogs at an astronomical rate).