Indietrochapter 2 study notes
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Atoms, Elements, and the Periodic Table
Definition and Nature of Elements
An element is a fundamental substance that cannot be chemically changed or broken down into simpler substances. There are 118 known elements, each with a unique name and symbol. Elements are the building blocks of all matter.
Examples of Element Symbols: H (Hydrogen), O (Oxygen), Fe (Iron), Na (Sodium), Cl (Chlorine), etc.
Abundance: Hydrogen and Helium are the most abundant in the universe, while Iron and Oxygen are most abundant on Earth.
The Periodic Table: Historical Development
The periodic table organizes elements based on their properties and atomic masses (now atomic numbers). Early chemists noticed similarities among groups of three elements (triads), such as Cl, Br, I and Li, Na, K. Dmitri Mendeleev arranged elements by increasing mass and grouped similar elements, leaving gaps for undiscovered elements and predicting their properties with remarkable accuracy.
Mendeleev's Predictions: He predicted the existence and properties of elements like gallium (eka-aluminum) and germanium (eka-silicon) before their discovery.
The Modern Periodic Table
Today, the periodic table consists of 7 periods (horizontal rows) and 18 groups (vertical columns). Elements are arranged by increasing atomic number. The table is a powerful tool for understanding chemical behavior and relationships among elements.
Classification of Elements
Groups and Periods
Elements are classified into groups (columns) and periods (rows) based on recurring chemical properties.
Group 1: Alkali Metals (Li, Na, K, Rb, Cs, Fr) – Soft, silvery, highly reactive with water, never found in elemental form in nature, combine with halogens in a 1:1 ratio.
Group 2: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra) – Shiny, silvery, less reactive than alkali metals, combine with halogens in a 1:2 ratio.
Group 17: Halogens (F, Cl, Br, I, At) – Colorful, corrosive nonmetals, highly reactive, exist in various physical states at room temperature.
Group 18: Noble Gases (He, Ne, Ar, Kr, Xe, Rn) – Colorless, inert gases, very low chemical reactivity.
Metals, Nonmetals, and Metalloids
Metals: Shiny, solid (except mercury), ductile, malleable, good conductors of heat and electricity, high melting and boiling points.
Nonmetals: Occur as solids, liquids, or gases; brittle; poor conductors; low melting and boiling points; variable colors.
Metalloids (Semimetals): Properties intermediate between metals and nonmetals; shiny but brittle; intermediate conductivity and melting/boiling points.
Properties of Matter
Extensive and Intensive Properties
Extensive properties depend on the amount of substance (e.g., mass, volume, length). Intensive properties are independent of sample size (e.g., density, melting point, temperature).
Physical and Chemical Properties
Physical properties: Observed without changing composition (e.g., color, melting point, electrical conductivity).
Chemical properties: Observed when a substance undergoes a chemical change (e.g., rusting, combustion, tarnishing).
Fundamental Laws of Chemistry
Law of Conservation of Mass
Antoine Lavoisier established that mass is neither created nor destroyed in chemical reactions. The total mass of reactants equals the total mass of products.
Law of Definite Proportions
Joseph Proust showed that a pure compound always contains the same elements in the same proportion by mass, regardless of the sample source.
Law of Multiple Proportions
John Dalton found that elements can combine in different ratios to form different compounds, and the mass ratios are small whole-number multiples of each other.
Compound | Formula | Mass Ratio (N:O) |
|---|---|---|
Nitric oxide | NO | 14:16 |
Nitrogen dioxide | NO2 | 14:32 |
Nitrous oxide | N2O | 28:16 |
Atomic Theory and Structure
Dalton's Atomic Theory
Dalton proposed that:
Elements are made of tiny particles called atoms.
Each element is characterized by the mass of its atoms.
Atoms of one element differ from those of other elements.
Atoms combine in fixed ratios to form compounds.
Chemical reactions rearrange atoms but do not change them.
Discovery of Subatomic Particles
Electrons
J.J. Thomson's cathode ray experiment demonstrated the existence of negatively charged particles (electrons) and allowed calculation of the charge-to-mass ratio:
Electron Charge and Mass
R.A. Millikan's oil drop experiment measured the charge of an electron (). Combining this with Thomson's ratio, the electron mass was calculated:
Protons and Neutrons
Rutherford's gold foil experiment revealed the existence of a small, dense, positively charged nucleus containing protons. Later, neutrons were discovered as neutral particles in the nucleus.
Summary of Subatomic Particles
Particle | Mass (g) | Mass (amu) | Charge (C) | Charge (e) |
|---|---|---|---|---|
Electron | 9.109 382 x 10−28 | 5.485 799 x 10−4 | −1.602 176 x 10−19 | −1 |
Proton | 1.672 622 x 10−24 | 1.007 276 | +1.602 176 x 10−19 | +1 |
Neutron | 1.674 927 x 10−24 | 1.008 665 | 0 | 0 |
Elements, Atomic Number, and Isotopes
The atomic number (Z) is the number of protons in the nucleus and defines the element. The mass number (A) is the sum of protons and neutrons. Isotopes are atoms of the same element with different numbers of neutrons but identical chemical properties.
Example: 75Se has 34 protons, 41 neutrons, and 34 electrons.
Additional info: These foundational concepts are essential for understanding chemical reactions, the structure of matter, and the organization of the periodic table.
