Chap_03A-Atoms_and_Elements_updated_10_1_2014

Chapter 3A: Atoms and Elements

2. Elements and Symbols

• Elements are the primary substances from which all other substances are built.

• Definition: Cannot be broken down into simpler substances.

• Naming Origins: Elements named after planets, mythological figures, minerals, colors, geographic locations, and famous individuals.

3. Element Symbols

• Symbol Format: Most elements have a one- or two-letter abbreviation.

• Capitalization Rule: Only the first letter is capitalized; if there is a second letter, it is lowercase.

  • Examples: Co (cobalt), C (carbon), O (oxygen).

• Historical Roots: Some symbols derive from Greek or Latin names.

  • Examples: Na (sodium from natrium), Fe (iron from ferrum).

4. Diatomic Elements

• Some elements form molecules consisting of two atoms.

  • Examples of diatomic elements: H2 (hydrogen), O2 (oxygen), N2 (nitrogen), F2 (fluorine), Cl2 (chlorine), Br2 (bromine), I2 (iodine).

5. Properties of Metals and Non-Metals

6. Metalloids

• Definition: Elements that exhibit properties of both metals and non-metals.

• Important Metalloids: Silicon (Si) and Germanium (Ge), crucial for technology, especially in computer chips.

7. The Periodic Table Overview

• Organization: Divided into periods (rows) and groups or families (columns).

• Elements in the same family possess similar properties.

• Trends:

  • Metallic character decreases across a period and increases down a group (e.g., Cs, Fr are most metallic; F is least metallic).

8. Groups and Periods

• Main-group Elements: Groups 1-2 and 13-18 are the main group or representative groups.

  • Alkali Metals (Group 1): Soft and very reactive, may react explosively.

  • Noble Gases (Group 18): Unreactive and used in applications like lighting.

  • Halogens (Group 17): Most reactive nonmetals, found in nature only as compounds.

  • Alkaline-earth Metals (Group 2): Less reactive than alkali metals.

  • Transition Metals: Metals located between main group elements.

9. Early Concepts of the Atom

• Atom Definition: The smallest particle of matter retaining its properties.

• Greek Philosopher Democritus (5th century B.C.) introduced the concept of atomos (indivisible particles).

10. Dalton’s Atomic Theory

• Year: 1808

• Model Description: Atom represented as a uniform density ball; referred to as "soccer ball" model.

• Three Parts of Dalton's Theory:

  • Each element is composed of tiny indestructible particles (atoms).

  • Atoms of a given element are similar but differ from other elements.

  • Atoms combine in simple, whole-number ratios to form compounds.

11. Discovery of the Electron

• Subatomic Particles: Smaller than atoms, called subatomic particles.

• J.J. Thomson's Experiment (1897): Used a cathode ray tube to discover electrons, demonstrating their negative charge.

• Atomic Model: Thomson proposed the "plum pudding" model with electrons embedded in a positively charged sphere.

12. Discovery of the Nucleus

• Ernest Rutherford’s Experiment (1910): Bombarded thin gold foil with alpha particles.

• Observations:

  • Majority passed through; some deflected at large angles; few turned back.

• Proposed a nuclear model of the atom: a small, dense nucleus surrounded by electrons.

13. The Modern Atom

• Described as a neutral spherical entity with a positively charged nucleus and negatively charged electrons.

• Protons (+) and neutrons (0) reside in the nucleus, while electrons (−) move rapidly around it.

14. Atomic Structure

• Subatomic Particles:

  • Proton: Charge +1, relative mass ~1800

  • Neutron: Charge 0, relative mass ~1800

  • Electron: Charge −1, relative mass 1

• Identity and Mass: Atomic number (Z)= number of protons; mass number (A) = number of protons + number of neutrons.

15. Isotopes

• Definition: Atoms of the same element with a different number of neutrons.

• Isotopes have the same atomic number but different mass numbers.

16. Atomic Mass and Isotopes

• Measured relative to a standard (carbon-12 atom) in atomic mass units (amu).

• Average atomic mass accounts for natural isotopic abundance.

• Example Calculation: # of protons, electrons, and neutrons in a chlorine atom. (P=17, E=17, N=18).

17. Isotope Examples

• To determine isotopes based on atomic mass and neutron difference.

• Common isotopes for Boron - distinguishing based on mass differences.

18. Conclusion

• The principles outlined in this chapter provide foundational knowledge of atomic theory and structure, including the development and understanding of elements, isotopes, and their properties.