Atoms and Elements: Structure, Laws, and the Periodic Table

Atoms and Elements

Introduction

This chapter introduces the foundational concepts of atoms and elements, their structure, the laws governing their behavior, and their organization in the periodic table. Understanding these principles is essential for further study in chemistry.

Imaging and Moving Individual Atoms

Atoms: The Fundamental Units of Matter

• Atom: The smallest identifiable unit of an element, typically with a radius of approximately $10^{-10}$ meters.

• Atoms are too small to be visualized with traditional optical microscopy.

Scanning Tunneling Microscopy (STM)

• STM was developed in 1981 and allows for imaging atoms on conductive surfaces.

• An atomically sharp tip is brought near a surface, and a tunneling current is maintained between the tip and the sample.

• As the tip moves across the surface, changes in current (due to the presence or absence of atoms) are detected, and the tip height is adjusted to maintain constant current.

• The resulting data allows for visualization of atomic arrangements, such as the hexagonal array of carbon atoms in graphite.

• Application: STM can also be used to manipulate individual atoms on surfaces.

Modern Atomic Theory and the Laws That Led to It

Key Laws Supporting Atomic Theory

• Law of Conservation of Mass: In a chemical reaction, matter is neither created nor destroyed. The total mass of reactants equals the total mass of products.

  • Example: $\mathrm{Na(s) + \frac{1}{2}Cl_2(g) \rightarrow NaCl(s)}$

  • If 7.7 g of Na and 11.9 g of Cl2 are combined, the total mass before and after the reaction remains the same (19.6 g).

• Law of Definite Proportions: All samples of a given compound have the same proportions by mass of their constituent elements.

  • Example: Decomposition of water always yields hydrogen and oxygen in a mass ratio of approximately 1:8.

• 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: Carbon and oxygen form CO and CO2; the ratio of oxygen masses that combine with 1 g of carbon in each compound is a small whole number ratio.

Dalton's Atomic Theory

• Elements are composed of indestructible particles called atoms.

• Atoms of a given element have the same mass and properties, distinguishing them from atoms of other elements.

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

• In chemical reactions, atoms are rearranged but not changed into other atoms.

Atomic Structure

Subatomic Particles

• Proton: Positively charged particle in the nucleus; mass ≈ 1.00727 u.

• Neutron: Neutral particle in the nucleus; mass ≈ 1.00866 u.

• Electron: Negatively charged particle outside the nucleus; mass ≈ 0.00054 u.

Particle	Symbol	Mass (u)	Charge (C)
Proton	p	1.00727	+1.60218 × 10-19
Neutron	n	1.00866	0
Electron	e-	0.00054	-1.60218 × 10-19

Atomic Number, Mass Number, and Isotopes

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

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

• Isotopes: Atoms of the same element with different numbers of neutrons (e.g., Ne-20, Ne-21, Ne-22).

• Some isotopes are radioactive and can emit particles or radiation (used in medical imaging and treatment).

Ions

• Anion: Atom with excess electrons (negatively charged), e.g., F-.

• Cation: Atom with fewer electrons than protons (positively charged), e.g., Na+.

• Ionization involves the gain or loss of electrons, not protons.

Atomic Mass

Average Atomic Mass and Isotopic Abundance

• The atomic mass listed in the periodic table is a weighted average of all naturally occurring isotopes of an element.

• Calculation: $\text{Atomic mass} = \sum (\text{fraction of isotope}_n) \times (\text{mass of isotope}_n)$

• Example (Chlorine):

  • Cl-35: 75.77% abundance, mass = 34.97 u

  • Cl-37: 24.23% abundance, mass = 36.97 u

  • Weighted average: $(0.7577 \times 34.97) + (0.2423 \times 36.97) = 35.45$ u

Mass Spectrometry

• Technique used to measure the masses and relative abundances of isotopes.

• Atoms are ionized, accelerated, and deflected by a magnetic field; the degree of deflection depends on mass-to-charge ratio ($m/z$).

• The resulting spectrum shows peaks corresponding to different isotopes; peak heights indicate relative abundance.

Molar Mass: Counting Atoms by Weighing Them

The Mole and Avogadro's Number

• Mole (mol): The amount of substance containing $6.022 \times 10^{23}$ entities (Avogadro's number).

• Used to relate atomic/molecular scale to macroscopic quantities.

• Molar Mass: The mass (in grams) of one mole of a substance; numerically equal to the atomic or molecular mass in u.

• Example: 12.01 g of carbon contains $6.022 \times 10^{23}$ atoms.

Conversions Involving Moles

• To convert between mass and moles: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• To convert between moles and number of particles: $\text{number of particles} = \text{moles} \times 6.022 \times 10^{23}$

The Periodic Table of the Elements

Organization and Classification

• Elements are arranged in order of increasing atomic number.

• Vertical columns are called groups or families; elements in the same group have similar chemical properties.

• Horizontal rows are called periods.

• Main categories:

  • Main-group elements (s- and p-blocks)

  • Transition elements (d-block)

  • Lanthanoids and Actinoids (f-block)

Groups of the Periodic Table

• Alkali Metals (Group 1): Very reactive, tend to lose 1 electron (1+ charge).

• Alkaline Earth Metals (Group 2): Reactive, tend to lose 2 electrons (2+ charge).

• Halogens (Group 17): Reactive nonmetals, exist as diatomic molecules (F2, Cl2, Br2), tend to gain 1 electron (1- charge).

• Noble Gases (Group 18): Very stable and unreactive.

Metals, Nonmetals, and Metalloids

• Metals: Good conductors of heat and electricity, malleable, ductile, tend to lose electrons.

• Nonmetals: Poor conductors, can be solid, liquid, or gas at room temperature, tend to gain electrons.

• Metalloids: Exhibit properties intermediate between metals and nonmetals; often semiconductors.

Category	Properties	Examples
Metals	Conductive, malleable, ductile, shiny	Fe, Cu, Na
Nonmetals	Poor conductors, brittle (if solid), various states	O, N, Cl
Metalloids	Intermediate properties, semiconductors	Si, B, As

Additional info: Periodic trends such as atomic radius, ionization energy, and electronegativity are discussed in later chapters, but the general classification and group properties are foundational for understanding chemical behavior.