Chapter 2 Lecture slides/notes

Atoms and Bonds

Introduction to Atoms and Ions

Atoms are the fundamental units of matter, composed of subatomic particles. When atoms lose or gain electrons, they become ions, which are charged particles essential for many biological processes.

• Atom: The smallest unit of an element, retaining its chemical properties.

• Ion: An atom or molecule with a net electric charge due to the loss or gain of electrons.

• Example: If an atom loses an electron, it becomes a positively charged ion (cation).

Chemical Elements and Compounds

Elements, Matter, and the Periodic Table

Elements are pure substances consisting of only one type of atom. Matter is anything that has mass and occupies space. The periodic table organizes elements based on their properties and atomic structure.

• Matter: Anything with mass and volume.

• Element: A substance that cannot be broken down into simpler substances by chemical means.

• Symbols: Element symbols are often derived from Latin or German names.

Essential and Trace Elements in Biology

Living organisms require certain elements in large (essential) or small (trace) amounts. These elements are critical for biological structure and function.

Element	Atomic Number	Percentage of Human Body Weight
Oxygen (O)	8	65.0
Carbon (C)	6	18.5
Hydrogen (H)	1	9.5
Nitrogen (N)	7	3.3
Calcium (Ca)	20	1.5
Phosphorus (P)	15	1.0
Potassium (K)	19	0.4
Sulfur (S)	16	0.3
Sodium (Na)	11	0.2
Chlorine (Cl)	17	0.2
Magnesium (Mg)	12	0.1

• Essential elements: Make up about 96% of body mass (O, C, H, N).

• Trace elements: Required in minute amounts but vital for health (e.g., Fe, I, Zn).

Emergent Properties of Compounds

When elements combine to form compounds, the resulting substances often have properties different from the individual elements.

• Emergent property: New characteristics that arise when components interact (e.g., sodium and chlorine form sodium chloride, which is edible salt).

Atoms and Molecules

Subatomic Particles

Atoms consist of a nucleus containing protons and neutrons, surrounded by electrons in orbitals.

• Protons: Positively charged particles in the nucleus.

• Neutrons: Neutral particles in the nucleus.

• Electrons: Negatively charged particles in orbitals around the nucleus.

Units of Measure in Atomic Structure

Atomic Number and Mass Number

The atomic number and mass number are fundamental for identifying and distinguishing atoms.

• Atomic Number (Z): Number of protons in the nucleus; also equals the number of electrons in a neutral atom.

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

Notation: AZElement (e.g., 126C)

Atomic Mass and the Dalton

Atomic mass is measured in Daltons (Da), where 1 Dalton ≈ 1.7 × 10-24 grams.

• Proton mass: 1 Da

• Neutron mass: 1 Da

• Electron mass: Negligible (1/2000 of a Dalton)

Example: Carbon-12 has 6 protons and 6 neutrons, so its atomic mass is 12 Da.

Isotopes

Stable and Unstable Isotopes

Isotopes are atoms of the same element with different numbers of neutrons. Some are stable, while others are radioactive (unstable).

• Stable isotopes: Do not change or decay over time (e.g., 12C, 13C).

• Unstable isotopes (radioisotopes): Decay spontaneously, emitting radiation (e.g., 14C decays to 14N).

Radioactive decay can convert one element to another by changing the number of protons.

Electron Shells and Valence Electrons

Electron Configuration

Electrons are arranged in shells around the nucleus. The distribution of electrons determines chemical reactivity.

• First shell: Holds up to 2 electrons.

• Second shell: Holds up to 8 electrons.

• Third shell: Holds up to 8 electrons (for main group elements).

Valence electrons are found in the outermost shell and are key to chemical bonding.

Inert Elements

Elements with full valence shells (e.g., noble gases) are chemically unreactive.

Chemical Bonds

Types of Chemical Bonds

Atoms form bonds to achieve stable electron configurations. The main types of bonds are:

• Polar covalent bonds

• Non-polar covalent bonds

• Ionic bonds

• Hydrogen bonds

• Van der Waals forces

All these interactions depend on electronegativity, the tendency of an atom to attract electrons.

Electronegativity

Electronegativity is a measure of an atom's ability to attract electrons in a bond.

• High electronegativity: Atom strongly attracts electrons (e.g., oxygen).

• Low electronegativity: Atom weakly attracts or loses electrons (e.g., sodium).

Non-Polar Covalent Bonds

Electrons are shared equally between atoms with similar electronegativity.

• Example: H2, O2, CH4

Polar Covalent Bonds

Electrons are shared unequally, resulting in partial charges on atoms.

• Example: H2O (oxygen is more electronegative than hydrogen)

Ionic Bonds

One atom donates an electron to another, creating oppositely charged ions that attract each other.

• Example: NaCl (sodium chloride)

Hydrogen Bonds

Weak attractions between the partial positive charge of a hydrogen atom and the partial negative charge of another atom (often oxygen or nitrogen).

• Example: Hydrogen bonding between water molecules

Van der Waals Forces

Weak, transient attractions due to temporary shifts in electron distribution.

• Individually weak, but collectively significant in large molecules.

Chemical Reactions and Equilibrium

Balanced Chemical Equations

Chemical equations must be balanced, meaning the number and type of atoms are the same on both sides of the reaction.

• Example: $2H_2 + O_2 \rightarrow 2H_2O$

Chemical Equilibrium

At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction. The concentrations of reactants and products remain constant.

• Example: The formation and breakdown of water can reach equilibrium under certain conditions.

Additional info: These foundational concepts are essential for understanding the chemical basis of life and are directly relevant to General Biology topics such as cell structure, metabolism, and genetics.