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

Chapter 4 – Atoms and Elements

Section 4.1: Elements and Symbols

Elements are the fundamental substances from which all matter is constructed. Each element is represented by a unique chemical symbol.

• Features of Elements:

  • Pure substances from which all other things are built.

  • Cannot be broken down into simpler substances by chemical means.

• Chemical Symbol Features:

  • Represent the names of the elements (e.g., C for carbon, S for sulfur).

  • Consist of one or two letters; the first letter is always capitalized, the second (if present) is lowercase (e.g., Au for gold, Al for aluminum).

Section 4.2: The Periodic Table

The periodic table organizes all known elements (currently 118) into a systematic arrangement based on their properties and atomic structure.

• Elements are organized into groups (vertical columns) and periods (horizontal rows).

• Elements are placed in order of increasing atomic mass (historically) and atomic number (modern table).

Groups and Periods

• Groups: Vertical columns containing elements with similar chemical properties.

• Periods: Horizontal rows numbered from 1 to 7.

Group Numbers and Names

• Group numbers are written at the top of each column.

• Representative elements: Groups 1A–8A (or 1, 2, 13–18).

• Transition elements: Groups 3B–8B, 1B, 2B (or 3–12).

• Alternative system: Groups numbered 1–18 from left to right.

Alkali & Alkaline Earth Metals

• Group 1A (1): Alkali Metals – lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs).

• Group 2A (2): Alkaline Earth Metals – beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), radium (Ra).

Halogens and Noble Gases

• Group 7A (17): Halogens – fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At), tennessine (Ts).

• Group 8A (18): Noble Gases – helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), radon (Rn), oganesson (Og).

Metals, Nonmetals, and Metalloids

• The heavy zigzag line on the periodic table separates metals (left), nonmetals (right), and metalloids (along the line).

Characteristics of Metals, Nonmetals, and Metalloids

Type	Location	Properties
Metals (except H)	Left side	Shiny, ductile, conduct heat/electricity, mostly solids (except Hg)
Nonmetals	Right side	Dull, brittle, poor conductors, often good insulators, low densities/melting points
Metalloids	Along zigzag line	Properties of both metals and nonmetals, semiconductors, used in electronics

Section 4.3: The Atom

An atom is the smallest particle of an element that retains its chemical properties.

• Atoms are composed of a nucleus (containing protons and neutrons) and electrons that occupy space around the nucleus.

Dalton’s Atomic Theory (1808)

• All matter is made of atoms.

• Atoms of the same element are identical; atoms of different elements are different.

• Atoms combine in fixed ratios to form compounds.

• Atoms are rearranged in chemical reactions but are not created or destroyed.

Additional info: Modern atomic theory recognizes isotopes and subatomic particles, which Dalton did not know.

Key Experiments

• 1897: J.J. Thomson discovered electrons (Plum Pudding Model).

• 1911: Ernest Rutherford discovered the nucleus via gold foil experiment.

Electromagnetic Radiation

• Energy that travels as waves; particles are called photons.

• Wavelength (λ): Distance between wave peaks (units: nm, μm, mm, etc.).

• High-energy radiation has short wavelengths; low-energy has long wavelengths.

Electromagnetic Spectrum

• Includes radio waves, microwaves, infrared, visible light, ultraviolet, X-rays, gamma rays.

Charged Particles in an Atom

• Protons: Positive charge (+1)

• Electrons: Negative charge (–1)

• Neutrons: No charge (neutral)

• Like charges repel; unlike charges attract.

Structure of the Atom

• Nucleus (center): Contains protons and neutrons; most of the atom’s mass.

• Electrons: Occupy space around the nucleus.

Mass of the Atom

• Measured in atomic mass units (amu): $1\ \text{amu} = \frac{1}{12}$ the mass of a carbon-12 atom ($1.66 \times 10^{-27}$ kg).

• 1 amu = 1 Dalton (Da).

• Electrons have negligible mass compared to protons and neutrons.

Particle	Symbol	Charge	Mass (amu)	Location
Proton	p or p+	+1	1.007	Nucleus
Neutron	n or n0	0	1.008	Nucleus
Electron	e–	–1	0.00055	Outside nucleus

Section 4.4: Atomic Number and Mass Number

• Atomic Number (Z): Number of protons in the nucleus; unique for each element; appears above the symbol in the periodic table.

• Mass Number (A): Total number of protons and neutrons in the nucleus; always a whole number; does not appear in the periodic table.

For neutral atoms: Number of protons = number of electrons.

• Cation: Atom with more protons than electrons (positive charge).

• Anion: Atom with more electrons than protons (negative charge).

Counting Protons, Electrons, and Neutrons

• Given atomic number (Z) and mass number (A):

  • Number of protons = Z

  • Number of electrons = Z (for neutral atom)

  • Number of neutrons = A – Z

Section 4.5: Isotopes and Atomic Mass

• Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons (different mass numbers).

• Isotopes are represented by their atomic symbols (e.g., $^{24}\text{Mg}$, $^{25}\text{Mg}$, $^{26}\text{Mg}$).

Average Atomic Mass

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

• Calculation:

$\text{Average atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

• Fractional abundance is the percent abundance divided by 100.

Isotope	Number of Protons	Number of Neutrons	Mass of Isotope (amu)	Percent Abundance (%)
$^{24}$Mg	12	12	23.99	78.70
$^{25}$Mg	12	13	24.99	10.13
$^{26}$Mg	12	14	25.98	11.17

Example Calculation: For magnesium, the average atomic mass is calculated as:

$\text{Average atomic mass} = (23.99 \times 0.7870) + (24.99 \times 0.1013) + (25.98 \times 0.1117) = 24.31\ \text{amu}$

Additional info: Weighted averages are used because not all isotopes are equally abundant in nature.