Chapter 3 Atomic Trends

Chapter 3: Atoms and Elements

Introduction to Elements

• Elements: Primary substances from which all other substances are built

  • Cannot be broken down into simpler substances

  • Building blocks of everything around us

Periodic Table and Element Symbols

• Elements have unique one or two-letter symbols derived from various sources:

  • Names after planets (e.g. Uranium from Uranus)

  • Mythical figures (e.g. Titanium from Titans)

  • Colors (e.g. Chlorine from 'chloros' meaning yellow)

  • Famous people (e.g. Einsteinium, Californium)

  • Elements often written with only the first letter capitalized, second letter in lowercase (e.g. Cobalt: Co, Carbon Monoxide: CO)

  • Some symbols are based on Latin or Greek names (e.g. Sodium: Na from 'natrium', Iron: Fe from 'ferrum')

Diatomic Elements

• Seven diatomic elements cannot exist as singular atoms:

  • Hydrogen (H2), Oxygen (O2), Nitrogen (N2), Fluorine (F2), Chlorine (Cl2), Bromine (Br2), Iodine (I2)

  • Identified on the periodic table in a '7' shape, except for Hydrogen

Structure of the Periodic Table

• Categories in the Periodic Table:

  • Metals: Shown in blue, generally good conductors

  • Nonmetals: Shown in yellow, opposite properties to metals

  • Metalloids: Shown in red, possess properties of both metals and nonmetals

    • Identified by a staircase line on the table (e.g. Boron, Silicon, Germanium)

• Properties of Metals and Nonmetals:

  • Metals: Solid, shiny, good conductors, malleable, ductile, lose electrons easily

  • Nonmetals: Gain or share electrons easily, generally brittle

  • Importance of losing (metals) and gaining (nonmetals) electrons

Periodic Trends

• Metallic Character: Decreases from left to right across the period

  • Most metallic: Cesium, Francium

  • Least metallic: Fluorine

• **Groups and Families: **

  • Periods: Rows (left to right)

  • Families: Columns (top to bottom)

    • Alkaline Metals (Group 1), Alkaline Earth Metals (Group 2), Transition Metals (Groups in between), Halogens (Group 17), Noble Gases (Group 18)

  • Noble gases are chemically inert and do not typically react

• Alkaline metals are very reactive and can explode with contact to water

Historical Development of Atomic Theory

• Early Concepts:

  • Democritus: Proposed matter is composed of small particles called 'atomos'

  • Dalton: Proposed elements made of identical atoms, refined the view of atoms

    • Introduced the "Soccer Ball Model" of uniform density

    • Atoms combine in whole number ratios to form compounds

Discovery of Subatomic Particles

• JJ Thomson:

  • Discovered the electron through cathode ray tube experiments

  • Proposed the "Plum Pudding Model" of the atom

• Ernest Rutherford:

  • Discovered the proton and nucleus via the Gold Foil Experiment

  • Proposed that atoms consist of a dense nucleus surrounded by electrons

Structure of the Modern Atom

• Composition of a neutral atom:

  • Nucleus: Center containing protons (+1 charge) and neutrons (0 charge)

  • Electrons: Surrounding the nucleus with a -1 charge

• Relative mass: Protons and neutrons are approximately 1,800 times heavier than electrons

Atomic Number and Mass Number

• Atomic Number (Z): Number of protons in an atom, unique to each element

• Mass Number (A): Sum of protons and neutrons

• In a neutral atom, the number of protons equals the number of electrons

• The mass number helps identify isotopes that have the same number of protons but different neutrons

Isotopes and Atomic Mass

• Isotopes: Atoms of the same element differing in the number of neutrons (e.g. Protium, Deuterium, Tritium of Hydrogen)

• Average atomic mass: Reflects the mass of an element’s naturally occurring isotopes based on their abundance

• Calculating Average Atomic Mass:

  • Multiply the mass of each isotope by its relative abundance

  • Add the calculated values to get the average atomic mass

Example of Calculating Atomic Information

• Chlorine Atom Calculation:

  • Atomic Mass: 35; Atomic Number: 17 (Protons)

  • Neutrons Calculation: 35 (Mass Number) - 17 (Protons) = 18 Neutrons

• Isotopes Example:

  • Isotopes Y and R: Same atomic number (85 protons) but different mass number (410 and 412) indicate Y & R are isotopes

  • Z and X: Different atomic numbers indicate they are not isotopes

Conclusion

• Understanding atomic structure, periodic trends, and isotopes is integral to the study of chemistry.