Chemistry: The Central Science - Chapter 2 Study Notes
Chemistry: The Central Science - Chapter 2 Study Notes
2.1 Atomic Theory of Matter
Historical Views:
Greek philosophers, notably Democritus, proposed the existence of indivisible particles termed “atomos” meaning uncuttable, which are the fundamental constituents of nature.
Formation of Atomic Theory:
In the early 1800s, John Dalton formalized atomic theory based on several laws:
Law of Constant Composition:
Discovered by Joseph Proust, this law states that compounds have a definite composition such that the relative number of atoms of each element is consistent across all samples of that compound.
Integral to Dalton's Atomic Theory (Postulate 4).
Law of Conservation of Mass:
Articulated by Antoine Lavoisier, this law asserts that the total mass of substances remains unchanged during a chemical reaction—mass is neither created nor destroyed.
Essential to Dalton's Atomic Theory (Postulate 3).
Law of Multiple Proportions:
This law states that when two elements combine to form more than one compound, the ratios of the masses of one element combined with a fixed mass of the other element are in the ratio of small whole numbers.
Example: Carbon monoxide (CO) vs. Carbon dioxide (CO₂).
Discovered as part of Dalton's development of atomic theory.
2.2 Discovery of Subatomic Particles
Emergence of Subatomic Theory:
Dalton initially viewed atoms as the smallest indivisible particles; however, subsequent discoveries revealed smaller components:
Cathode Rays:
J.J. Thomson discovered electrons as streams of negatively charged particles in 1897, demonstrating that atoms were not indivisible.
Radioactivity:
Henry Becquerel first observed spontaneous emission of radiation, later studied by Marie and Pierre Curie, revealing additional subatomic phenomena.
Structure of the Atom:
The atom comprises a nucleus (containing protons and neutrons) and electrons that orbit around it.
Measurements and Discovery of Electrons:
Thomson’s discoveries led to the calculation of the electron's charge-to-mass ratio.
Millikan Oil-Drop Experiment:
Conducted by Robert Millikan in 1909, determining the electron charge as approximately 1.602 imes 10^{-19} ext{ coulombs} and confirming its mass.
Types of Radiation:
Rutherford's Discoveries:
Conducted experiments demonstrating three radiation types:
Alpha particles (positively charged), beta particles (negatively charged), and gamma rays (neutral).
2.3 Modern View of Atomic Structure
Rutherford's Nuclear Model:
Postulated a small, dense, positively charged nucleus surrounded by electrons orbiting in a vast empty space,
Atoms are primarily empty space, with their size measured in picometers (1-5 ext{ Å}, or 100-500 ext{ pm}).
Subatomic Particle Characteristics:
Comparison of subatomic particles:
Particle
Charge
Mass (amu)
Proton
+1
1.0073
Neutron
0
1.0087
Electron
-1
5.486 imes 10^{-4}
2.4 Atomic Number
Definition of Atomic Number:
The atomic number is defined as the number of protons in an atom's nucleus, indicating an element's identity. Since atoms are electrically neutral, the number of electrons equals the number of protons.
Mass Number:
The mass number combines the total number of protons and neutrons in a nucleus, typically denoted as a superscript.
2.5 Isotopes
Definition of Isotopes:
Isotopes are variants of the same chemical element that contain equal numbers of protons but different numbers of neutrons, leading to variations in atomic mass.
Example Table for Carbon Isotopes:
Symbol
Protons
Electrons
Neutrons
C-11
6
6
5
C-12
6
6
6
C-13
6
6
7
C-14
6
6
8
2.6 Atomic Mass Unit (amu)
Definition and Importance:
An atomic mass unit (amu) is a unit of mass used to measure atomic and molecular weights, defined such that the mass of carbon-12 is exactly 12 amu.
This is crucial for specifying relative atomic weights accurately.
2.7 Atomic Weight
Definition of Atomic Weight:
Atomic weight represents the average of the isotopic masses of an element, weighted by their natural abundances.
Measurement Techniques:
Measurement of atomic and molecular weights can often be performed using a mass spectrometer, providing detailed isotopic distribution data.
2.8 Periodic Table
Structure of the Periodic Table:
Elements arranged according to increasing atomic numbers enable systematic study. Accessible periodic tables can be located at provided URLs.
Groups and Periods:
Definition of Groups: Vertical columns with elements sharing similar chemical properties.
Definition of Periods: Horizontal rows indicative of periodic trends in chemical properties, with reactivity showing recurrent patterns.
2.9 Types of Compounds
Molecular Compounds:
Defined as compounds composed of molecules typically containing nonmetals.
Diatomic Molecules:
Seven elements exist as diatomic molecules: H, N, O, F, Cl, Br, I. These can be identified in the periodic table.
2.10 Ions and Ionic Compounds
Ion Formation:
Ions are formed when atoms gain or lose electrons:
Cations: Positively charged ions formed by losing electrons.
Anions: Negatively charged ions formed by gaining electrons.
Creation of Ionic Compounds:
Ionic compounds (e.g., NaCl) are formed by the transfer of electrons between metals and nonmetals, culminating in the attraction of oppositely charged ions.
2.11 Naming Inorganic Compounds
Chemical Nomenclature:
The naming system for ionic compounds necessitates the memorization of common ions. Standard rules apply consistently across different types of compounds (Ionic, Acids, Binary molecular).
Common Ions Table
Cations:
Charge
Formula
Name
1+
H+
Hydrogen ion
1+
Li+
Lithium ion
1+
Na+
Sodium ion
2+
Ca2+
Calcium ion
2+
Zn2+
Zinc ion
Anions:
Charge
Formula
Name
1-
F-
Fluoride ion
2-
O2-
Oxide ion
3-
N3-
Nitride ion
Summary of Inorganic Compounds Naming
The cation precedes the anion in naming. An cation's charge may be expressed in Roman numerals if it has multiple oxidation states.
2.12 Organic Compounds
Overview of Organic Chemistry:
Organic compounds focus primarily on carbon-containing substances, with nomenclature specific to carbon-based functional groups.
Nomenclature of Alkanes:
Alkanes are defined as hydrocarbons featuring only single bonds, suffixed with -ane (e.g., methane, ethane).
Isomers Explanation:
Isomers are substances with the same molecular formula but different structural configurations (e.g., 1-Propanol and 2-Propanol).
Nomenclature of Alcohols:
Produced by replacing one hydrogen in an alkane with a hydroxyl group (-OH), and typically ending the name with -ol.