Atoms and Elements

Definition of Atoms and Elements

Atoms are the fundamental building blocks of matter, consisting of a nucleus surrounded by electrons. Elements are pure substances composed of only one type of atom, each defined by its unique number of protons.

• Atom: The smallest unit of an element that retains its chemical properties.

• Element: A substance that cannot be broken down into simpler substances by chemical means. Each element is identified by its atomic number (number of protons).

• Example: Oxygen (O) is an element; each oxygen atom has 8 protons.

Modern Atomic Theory

Dalton's Atomic Theory

John Dalton proposed the first modern atomic theory, which laid the foundation for our understanding of matter.

• All matter is composed of indivisible atoms.

• Atoms of the same element are identical in mass and properties.

• Atoms of different elements differ in mass and properties.

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

• Chemical reactions involve the rearrangement of atoms; atoms are not created or destroyed.

• Example: Water (H2O) is formed by combining two hydrogen atoms and one oxygen atom.

Fundamental Laws of Chemistry

Law of Conservation of Mass

This law states that mass is neither created nor destroyed in a chemical reaction.

• Key Point: The total mass of reactants equals the total mass of products.

• Equation:

• Example: Burning hydrogen in oxygen produces water, but the total mass remains constant.

Law of Definite Proportions (Constant Composition)

A chemical compound always contains the same elements in the same proportion by mass.

• Key Point: Water always contains 11.2% hydrogen and 88.8% oxygen by mass.

• Equation:

Law of Multiple Proportions

When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

• Example: CO and CO2 both contain carbon and oxygen, but the ratio of oxygen in CO2 to CO is 2:1.

Discovery of Subatomic Particles

Discovery of the Electron

Electrons were discovered by J.J. Thomson using the cathode ray tube experiment.

• Cathode Ray Tube: A vacuum tube that emits a beam of electrons (cathode rays) when a voltage is applied.

• Thomson's Plum Pudding Model: Proposed that atoms are spheres of positive charge with negatively charged electrons embedded within.

Rutherford's Nuclear Theory

Ernest Rutherford's gold foil experiment led to the nuclear model of the atom.

• Most of the atom's mass and all of its positive charge are concentrated in a small nucleus.

• Electrons move around the nucleus in the empty space of the atom.

• Example: The atom is mostly empty space, with a dense, positively charged nucleus at the center.

Subatomic Particles

Protons, Neutrons, and Electrons

Atoms are composed of three main subatomic particles:

• Proton (p+): Positively charged particle found in the nucleus; mass ≈ 1 amu.

• Neutron (n0): Neutral particle found in the nucleus; mass ≈ 1 amu.

• Electron (e-): Negatively charged particle found outside the nucleus; mass ≈ 0.0005 amu.

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Total number of protons and neutrons in the nucleus.

• Equation: (where N = number of neutrons)

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

• Notation: , where X is the element symbol, A is the mass number, and Z is the atomic number.

Ions: Cations and Anions

Ions are charged atoms or molecules formed by the loss or gain of electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Example: Na+ (sodium cation), Cl- (chloride anion).

The Periodic Table

Classification of Elements

The periodic table organizes elements by increasing atomic number and similar chemical properties.

• Groups (Columns): Elements with similar properties.

• Periods (Rows): Elements with increasing atomic number.

• Metals: Good conductors, malleable, ductile, shiny.

• Nonmetals: Poor conductors, brittle, dull.

• Metalloids: Properties intermediate between metals and nonmetals.

• Alkali Metals (Group 1): Highly reactive metals.

• Alkaline Earth Metals (Group 2): Reactive metals.

• Halogens (Group 17): Very reactive nonmetals.

• Noble Gases (Group 18): Inert, unreactive gases.

Ions and the Periodic Table

Elements tend to form ions to achieve a stable electron configuration.

• Metals tend to lose electrons and form cations.

• Nonmetals tend to gain electrons and form anions.

• Example: Na (metal) forms Na+; Cl (nonmetal) forms Cl-.

Atomic Mass and Molar Mass

Atomic Mass

The atomic mass of an element is the weighted average mass of its naturally occurring isotopes, measured in atomic mass units (amu).

• Example: The atomic mass of chlorine is approximately 35.45 amu.

Molar Mass and the Mole

The molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol). The mole is a counting unit used in chemistry.

• Avogadro's Number: particles per mole.

• Conversion:

• Example: 18.0 g of water (H2O) is 1 mole.

Conversions Between Mass, Moles, and Atoms

Conversions are essential for quantitative chemical calculations.

• Mass to Moles:

• Moles to Atoms:

• Atoms to Moles:

• Moles to Mass:

• Example: How many atoms are in 2.00 g of hydrogen? Step 1: Find moles: Step 2: Find atoms: atoms.

Summary Table: Subatomic Particles