Atoms, Chemical Bonds, and Water: Review Notes (Chapters 1-3)

Atoms and Atomic Structure

Basic Parts of an Atom

Atoms are the fundamental units of matter, composed of three main subatomic particles.

• Proton: Positively charged particle found in the nucleus.

• Neutron: Neutral particle (no charge) also located in the nucleus.

• Electron: Negatively charged particle that orbits the nucleus in electron shells.

Example: A sodium (Na) atom has 11 protons, 11 electrons, and (typically) 12 neutrons.

Electron Shells and Orbitals

Electrons are arranged in energy levels or shells around the nucleus. Each shell contains one or more orbitals, which are regions where electrons are likely to be found.

• First shell: Holds up to 2 electrons (1s orbital).

• Second shell: Holds up to 8 electrons (2s and 2p orbitals).

• 2s orbital: Spherical in shape, holds up to 2 electrons.

Example: Sodium (Na) has 1 electron in its valence (outermost) shell.

Electronegativity and Chemical Bonds

Electronegativity

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

• Higher electronegativity means a stronger pull on shared electrons.

• Electronegativity differences between atoms determine the type of bond formed.

Example: Oxygen is more electronegative than hydrogen, leading to polar covalent bonds in water.

Trends in Electronegativity

Electronegativity increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Most electronegative elements: Fluorine (highest), then oxygen, nitrogen, etc.

• Group review question: Of C, Si, K, S – S (sulfur) is the most electronegative.

• Least electronegative: Of N, Se, K, P – K (potassium) is the least electronegative.

Types of Chemical Bonds

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

• Ionic bond: Transfer of electrons from one atom to another, forming ions (e.g., NaCl).

• Covalent bond: Sharing of electron pairs between atoms.

• Polar covalent bond: Unequal sharing of electrons due to differences in electronegativity (e.g., H2O).

• Non-polar covalent bond: Equal sharing of electrons (e.g., O2).

• Hydrogen bond: Weak attraction between a hydrogen atom (covalently bonded to an electronegative atom) and another electronegative atom (e.g., between water molecules).

Strength of Bonds: Covalent bonds are the strongest, followed by ionic, then hydrogen bonds.

Bond Examples and Lewis Structures

Lewis structures show the arrangement of electrons in molecules. Shared pairs (bonds) and lone pairs (unshared electrons) are indicated.

• Shared electrons: Form covalent bonds.

• Paired electrons: Non-bonding (lone pairs).

• Partial charge: Occurs in polar covalent bonds, where electrons are unequally shared.

Example Table: Types of Bonds in Molecules

Bond Type	Electron Behavior	Example
Ionic	Transfer	NaCl
Polar Covalent	Unequal sharing	H2O
Non-polar Covalent	Equal sharing	O2
Hydrogen	Attraction between molecules	Between H2O molecules

Water: Structure and Properties

Water Molecules and Hydrogen Bonding

Water molecules are polar, with partial positive charges on hydrogen atoms and a partial negative charge on oxygen. This allows them to form hydrogen bonds with each other.

• Covalent bonds: Hold the H and O atoms together within a water molecule (polar covalent).

• Hydrogen bonds: Form between the hydrogen of one water molecule and the oxygen of another.

Example: In a cluster of water molecules, each molecule can form up to four hydrogen bonds.

Cohesion and Adhesion

Cohesion is the attraction between molecules of the same substance (e.g., water to water), while adhesion is the attraction between different substances (e.g., water to glass).

• Cohesion: Responsible for surface tension in water.

• Adhesion: Helps water climb up plant vessels (capillary action).

Density of Water and Ice

Water is most dense at 4°C. As water freezes, it forms a crystalline structure maintained by hydrogen bonds, making ice less dense than liquid water.

• Density decreases as water freezes: Ice floats on liquid water.

• Biological significance: Aquatic life can survive under ice in winter.

Example Table: Density of Water at Different Temperatures

Temperature (°C)	Density (g/cm3)
0 (ice)	0.920
4 (liquid)	1.000
10 (liquid)	0.998

Acids, Bases, and the pH Scale

pH Scale

The pH scale measures the concentration of hydrogen ions [H+] in a solution. It ranges from 0 (most acidic) to 14 (most basic).

• Acidic solutions: pH < 7, higher [H+]

• Neutral solution: pH = 7

• Basic (alkaline) solutions: pH > 7, higher [OH-]

Formula:

$\text{pH} = -\log_{10}[\text{H}^+]$

Example: Stomach acid has a pH of about 2 (very acidic), while household bleach is basic (pH ~12).

Applications and Biological Relevance

• Enzyme activity is highly dependent on pH.

• Organisms maintain homeostasis by regulating internal pH.

Additional info: Some context and examples were inferred to provide a complete, self-contained study guide based on the review questions and images.