Atoms, Elements, and Periodic Trends: Study Notes for Introductory Chemistry

Chapter 4: Atoms and Elements

Historical Development of Atomic Theory

The understanding of the atom has evolved through the contributions of several key scientists. Their discoveries have shaped modern chemistry.

• Democritus (c. 460 BCE): Proposed that all matter is made up of indivisible particles called atoms.

• Dalton: Developed the atomic theory, stating that atoms are the fundamental units of matter, participate in chemical reactions, and combine in whole-number ratios to form compounds.

• J. J. Thomson: Discovered the electron using the cathode ray tube experiment and proposed the Plum Pudding Model of the atom.

• Ernest Rutherford: Conducted the gold foil experiment, discovering the nucleus and proposing a nuclear model of the atom.

The Atom and Subatomic Particles

Atoms are composed of three main subatomic particles: protons, neutrons, and electrons. These particles determine the identity and properties of each element.

• Proton: Positively charged particle found in the nucleus.

• Neutron: Neutral particle found in the nucleus.

• Electron: Negatively charged particle found outside the nucleus.

• 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 (and thus different mass numbers).

Isotopes and Atomic Mass

Isotopes are variants of a particular chemical element that have the same number of protons but different numbers of neutrons. The atomic mass of an element is a weighted average of the masses of its naturally occurring isotopes.

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

• Average atomic mass: Calculated using the relative abundance and mass of each isotope.

Example: Chlorine has two main isotopes: 35Cl and 37Cl. The average atomic mass is calculated based on their natural abundances.

Mass Spectrometry

Mass spectrometry is a technique used to determine the masses and relative abundances of isotopes in a sample. It helps in identifying isotopes and calculating average atomic mass.

The Periodic Table

The periodic table organizes elements based on their atomic number and recurring chemical properties. Elements are grouped into periods (rows) and groups (columns).

• Groups: Vertical columns; elements in the same group have similar chemical properties.

• Periods: Horizontal rows.

• Metals, Nonmetals, Metalloids: Elements are classified based on their physical and chemical properties.

• Special groups: Alkali metals, alkaline earth metals, transition metals, halogens, noble gases.

Periodic Table and Ion Charges

The periodic table can be used to determine the typical charges of ions formed by elements.

• Metals: Tend to lose electrons and form cations (positive ions).

• Nonmetals: Tend to gain electrons and form anions (negative ions).

Chapter 9: Electron Configuration and Periodic Trends

Light and Energy

Light exhibits both wave-like and particle-like properties. The energy of light is quantized and can be described using specific equations.

• Wavelength (λ): Distance between two crests, measured in nanometers (nm).

• Frequency (ν): Number of waves that pass a point per second, measured in hertz (Hz).

• Energy of a photon:

$E=h\dfrac{c}{\lambda}$

• h: Planck's constant ($6.626 \times 10^{-34}$ J·s)

• c: Speed of light ($3.00 \times 10^8$ m/s)

Example: Calculate the energy of a photon with a wavelength of 500 nm.

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals. The modern model uses quantum numbers and the concept of orbitals.

• Principal quantum number (n): Indicates the main energy level.

• Orbital: Region of space where there is a high probability of finding an electron.

• Electron configuration notation: Shows the distribution of electrons among the atom's orbitals (e.g., 1s2 2s2 2p6).

• Orbital diagrams: Visual representations using arrows to indicate electron spins.

Bohr Model vs. Quantum Mechanical Model

The Bohr model describes electrons in fixed orbits around the nucleus, while the quantum mechanical model describes electrons in terms of probability distributions (orbitals).

• Bohr "orbit": Electrons travel in fixed paths at set distances from the nucleus.

• Quantum "orbital": Electrons exist in regions of space with a certain probability of being found.

Periodic Trends

Periodic trends are patterns observed in the properties of elements as you move across periods or down groups in the periodic table.

• Atomic radius: Generally decreases across a period and increases down a group.

• Ionization energy: Energy required to remove an electron from an atom; increases across a period and decreases down a group.

• Electron affinity: Tendency of an atom to accept an electron; generally becomes more negative across a period.

• Metallic character: Increases down a group and decreases across a period.

Electron Configurations and Ions

Electron configurations can be used to predict the number of unpaired electrons and the charges of ions formed by elements.

• Cations: Formed by loss of electrons; electron configuration is that of the nearest noble gas.

• Anions: Formed by gain of electrons; electron configuration is that of the nearest noble gas.

• Isoelectronic series: A group of ions and/or atoms that have the same number of electrons.

Trends in Ion Size

The size of ions depends on their charge and the number of electrons relative to the parent atom.

• Cations: Smaller than their parent atoms due to loss of electrons.

• Anions: Larger than their parent atoms due to gain of electrons.

• Isoelectronic series: Among isoelectronic species, the one with the most protons is the smallest.

Summary Table: Periodic Trends

Trend	Across a Period (→)	Down a Group (↓)
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electron Affinity	Becomes more negative	Becomes less negative
Metallic Character	Decreases	Increases

Key Equations

• Energy of a photon:

$E=h\dfrac{c}{\lambda}$

• Relationship between wavelength and frequency:

$c = \lambda \nu$

Example: Electron Configuration

Write the electron configuration for oxygen (O, atomic number 8):

• 1s2 2s2 2p4

Example: Predicting Ion Size

Compare the sizes of Na+, Mg2+, and F- (all isoelectronic with Ne): Mg2+ is the smallest, followed by Na+, then F-.