Atoms and Elements

Definition and Properties of Elements

Elements are the simplest substances that cannot be broken down by chemical means. Each element is represented in the Periodic Table, which provides information about the element's composition and properties.

• Examples: The four most common elements in living organisms by body weight are oxygen (O), carbon (C), hydrogen (H), and nitrogen (N). These elements make up about 96.3% of body weight.

• Each element is made up of only one kind of atom.

Structure of Atoms

Atoms are the building blocks of matter. They are extremely small and consist of three main subatomic particles: protons, neutrons, and electrons.

*1 atomic mass unit (amu) = 1.66 × 10-24 grams.

• Protons determine the atomic number and identity of the element.

• Neutrons add mass and can vary in number, resulting in isotopes.

• Electrons are involved in chemical bonding and reactions.

Atomic Number and Mass Number

• Atomic number (Z): Number of protons in the nucleus. Determines the element.

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

• For a neutral atom: Number of protons = Number of electrons.

Example: Carbon has an atomic number of 6 (6 protons). The most common isotope, carbon-12, has 6 neutrons (6 protons + 6 neutrons = mass number 12).

Isotopes

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

• Example: Carbon-12 (6 protons, 6 neutrons), Carbon-13 (6 protons, 7 neutrons), Carbon-14 (6 protons, 8 neutrons).

• Some isotopes, like carbon-14, are radioactive and decay spontaneously, emitting energy.

Electron Arrangement and Energy Levels

Electron Shells and Valence Electrons

Electrons are arranged in specific energy levels or shells around the nucleus. Each shell can hold a certain maximum number of electrons.

• Electrons fill the lowest energy levels first.

• Valence electrons are electrons in the outermost shell. They determine chemical reactivity.

• Atoms with a full valence shell are stable and are called inert gases (e.g., noble gases).

Ions and Ionic Bonding

Formation of Ions

An ion is an atom that has gained or lost electrons, resulting in a net charge.

• Cation: Positively charged ion (lost electrons). Example: Na+, K+, Mg2+, Ca2+.

• Anion: Negatively charged ion (gained electrons). Example: Cl-, F-.

Ionic Bonds

Ionic bonds form when electrons are transferred from one atom to another, resulting in oppositely charged ions that attract each other.

• Example: Sodium chloride (NaCl) forms when Na donates an electron to Cl, creating Na+ and Cl- ions that bond together.

Covalent Bonding

Definition and Types

Covalent bonds form when two atoms share one or more pairs of electrons. Covalent bonds can be single, double, or triple bonds depending on the number of shared electron pairs.

• Single covalent bond: One pair of electrons shared (e.g., H2).

• Double covalent bond: Two pairs of electrons shared (e.g., O2).

• Triple covalent bond: Three pairs of electrons shared (e.g., N2).

Polar and Nonpolar Covalent Bonds

The sharing of electrons in covalent bonds can be equal or unequal, depending on the atoms' electronegativity (the tendency of an atom to attract electrons).

• Nonpolar covalent bond: Electrons are shared equally (usually between atoms of the same element). Example: O2 gas.

• Polar covalent bond: Electrons are shared unequally (between atoms of different elements with different electronegativities). Example: H2O (water).

Molecules and Compounds

Definitions

• Molecule: Two or more atoms held together by covalent bonds. Example: O2, H2O.

• Compound: A substance containing two or more different elements chemically bonded together. Example: NaCl, H2O.

• All compounds are molecules, but not all molecules are compounds (e.g., O2 is a molecule but not a compound).

Examples

• Oxygen gas (O2): Molecule, not a compound (contains only one element).

• Methane (CH4): Both a molecule and a compound (contains carbon and hydrogen).

• Sodium chloride (NaCl): Compound, not a molecule (ionic bond, not covalent).

Electronegativity and Bond Prediction

Electronegativity

Electronegativity is a measure of how strongly an atom attracts electrons in a chemical bond. Differences in electronegativity between atoms determine the type of bond formed.

• Large difference: Ionic bond

• Small or no difference: Nonpolar covalent bond

• Intermediate difference: Polar covalent bond

Bond Prediction Rules

• Elements on opposite sides of the periodic table usually form ionic bonds.

• Elements on the right side of the periodic table usually form covalent bonds.

• Hydrogen (H) almost always forms covalent bonds.

Summary Table: Key Terms and Examples

Key Formulas and Equations

• Atomic number:

• Mass number:

• Number of neutrons:

Additional info:

• For most introductory biology courses, only the first 18 elements (up to argon) are considered for electron shell structure.

• The actual arrangement of electrons in atoms is more complex and involves subshells and orbitals, but this is beyond the scope of this summary.