Fundamental Concepts in GOB Chemistry: Atoms, Elements, and the Periodic Table

Atoms, Elements, and Molecules

Introduction to Biochemistry and Chemistry Systems

Understanding chemistry begins with breaking down systems into their component parts and studying their properties. In biochemistry, this involves analyzing molecules such as proteins, nucleic acids, carbohydrates, and lipids, and understanding their roles in metabolism and energy flow.

• System Analysis: Deconstruct to component parts, study isolated components, and reconstruct to understand the whole.

• Components of Life: Proteins, nucleic acids, carbohydrates, lipids.

• Energetics of Life: Metabolism and free energy.

Elements Found in Cells

Elements are the building blocks of biological molecules. Their structure determines their function in living systems.

• Structure vs. Function: The arrangement of atoms in molecules determines their biological role.

• Periodic Table: Elements are organized by atomic number and grouped by similar properties.

Definitions of Key Terms

• Matter: Physical material of the universe; anything that has mass and occupies space (volume).

• Mass: Amount of matter present.

• Pure Substance: Uniform chemical composition.

• Mixture: Variable composition; can be physically separated.

• Molecule: Smallest unit of pure substance that retains properties; made up of atoms.

• Atom: Smallest particle of matter (element).

• Element: Homogeneous, pure substance; only one kind of atom.

• Compound: Homogeneous, pure substance; multiple kinds of atoms.

Representations of Molecules

Molecules can be represented in several ways to illustrate their structure and bonding.

• Molecular Formula: Shows the elements and number of atoms (e.g., H2O).

• Line Structure: Shows bonds between atoms.

• Ball-and-Stick Model: Each element is shown as a colored ball connected by sticks (bonds).

• Space-Filling Model: Shows the relative size and position of atoms in a molecule.

Subatomic Particles and Atomic Structure

Subatomic Particles

Atoms are made of subatomic particles: protons, neutrons, and electrons. Their arrangement determines the properties of elements and molecules.

• Proton: Positively charged particle in the nucleus.

• Neutron: Neutral particle in the nucleus.

• Electron: Negatively charged particle in orbitals around the nucleus.

Particle	Symbol	Charge	Mass (amu)	Location
Proton	p	+1	1.007	Nucleus
Neutron	n	0	1.008	Nucleus
Electron	e	-1	0.00055	Orbitals

Atomic Number and Mass Number

• Atomic Number (Z): Number of protons in an element.

• Mass Number (A): Number of protons plus neutrons.

• Isotopes: Atoms with the same number of protons but different numbers of neutrons.

• Atomic Weight: Average mass of an element, based on isotope abundance.

Formula:

$A = Z + N$

Where $A$ is the mass number, $Z$ is the atomic number, and $N$ is the number of neutrons.

Molecular Weight

The molecular weight is the sum of the atomic weights for all atoms in a molecule.

• Water (H2O): $2 \times 1 + 16 = 18$

• Glucose (C6H12O6): $6 \times 12 + 12 \times 1 + 6 \times 16 = 180$

The Periodic Table and Electronic Structure

Periodic Table Structure

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

• Horizontal rows are called "periods".

• Vertical columns are called "groups" or "families".

• Elements in the same group have similar properties.

Electronic Structure of the Atom

Each element has a specific atomic number and a unique arrangement of electrons.

• Number of protons = atomic number.

• Number of electrons = number of protons (in a neutral atom).

• Electrons are located in energy shells outside the nucleus.

Energy Shells and Electron Location

Electrons occupy energy shells, and their location determines the energy and stability of the atom.

• Shell (principal quantum number n): $n = 1, 2, 3, ...$

• Subshell (shape): s, p, d, f

• Orbital: division of subshell (region of space)

• Spin: up or down

Each electron must have a different set of quantum numbers.

Valence Electrons

• Valence shell: Outermost electron shell of an element.

• Valence electron: Electron in the valence shell.

• Valence electrons determine chemical reactivity and bonding.

Periodic Trends and Properties

Metallic Character

Metallic elements conduct heat and electricity, form wires and sheets, and have metallic luster. The most metallic elements are at the bottom left of the periodic table.

Periodic Property: Atomic Size

• Atomic size increases down a group (more shells).

• Atomic size decreases across a period (increased nuclear charge pulls electrons closer).

• The largest atoms are at the bottom left; the smallest are at the top right.

Periodic Property: Ionization Energy

• Ionization Energy: Energy required to remove an electron ($E_I$).

• Ionization energy increases across a period and decreases down a group.

• Highest ionization energies are at the top right; lowest at the bottom left.

Formula:

$E_I = E_{final} - E_{initial}$

Noble Gases, Octet Rule, and Bonding

Noble Gases

• Colorless gases, generally inert and unreactive.

• 8 valence electrons (except He, which has 2).

• Full valence shell is stable; noble gases do not react with other elements.

Octet Rule

Atoms tend to gain, lose, or share electrons to achieve a full valence shell of 8 electrons, which is most stable.

Lewis Symbols and Bonding

• Lewis symbols show the number of valence electrons as dots around the element symbol.

• Atoms can gain or lose (transfer) electrons (ionic bonding) or share electrons (covalent bonding) to achieve a stable octet.

Basic Ionic Bonding Example

• Na transfers an electron to Cl: $\text{Na} \rightarrow \text{Na}^+ + e^-$

• Cl gains electron: $\text{Cl} + e^- \rightarrow \text{Cl}^-$

• Na+ and Cl- form NaCl (ionic compound).

Basic Covalent Bonding Example

• Each atom needs one more electron and has one unpaired electron in a valence shell.

• Electrons are shared between atoms to form a covalent bond.

Ions and Ionic Compounds

Definitions

• Neutral atom: Equal numbers of protons and electrons.

• Simple ion: Unequal numbers of protons and electrons.

• Cation: Fewer electrons than protons; positively charged.

• Anion: More electrons than protons; negatively charged.

Ionization and Ion Formation

Atoms can lose or gain electrons to form ions. Metals tend to lose electrons (form cations), while nonmetals tend to gain electrons (form anions).

Polyatomic Ions

Polyatomic ions are covalently bonded atoms with an overall molecular charge.

Ion	Name
NH4+	Ammonium
OH-	Hydroxide
HCO3-	Bicarbonate
CO32-	Carbonate
SO42-	Sulfate
PO43-	Phosphate

Ionic Bonds and Compounds

• Electrostatic attraction between oppositely charged ions forms ionic bonds.

• Binary compounds are formed from positive (cation) and negative (anion) ions.

• Subscripts represent the lowest possible ratio of ions.

Summary of Key Concepts

• An atom is made of subatomic particles (protons, electrons, neutrons); its properties are defined primarily by its subatomic structure.

• The atomic number is the number of protons and defines the element.

• Isotopes are atoms with the same number of protons but different numbers of neutrons.

• The atomic weight of an element depends on the relative abundance of its isotopes.

• The periodic table is organized by the number of valence electrons, giving rise to periodic trends: metallic character, atomic size, and ionization energy.

• Valence electrons determine chemical and physical properties.

• Ionic and covalent bonds are formed when atoms transfer or share electrons to achieve stable electron configurations.

• Polyatomic ions are important in biological and chemical systems.