Ch 2 Chemistry of Life1

Chemistry of Life

Parts of Atoms

Chapter 2.1 Learning Goals

• Describe the mass and charge of protons, electrons, and neutrons.

• Identify the atomic number and atomic mass of an atom based on the number of protons and neutrons.

What is Matter?

• Living organisms are made up of matter which occupies space and has mass.

• Matter consists of chemical elements, defined as substances that cannot be broken down by ordinary chemical means.

Elements

• Total of 118 elements; fewer than 30 are found in living cells.

• Unique properties for each element.

• Four elements make up about 96% of the mass of most living organisms:

  • Oxygen (O): 65%

  • Carbon (C): 18.5%

  • Hydrogen (H): 9.5%

  • Nitrogen (N): 3.3%

• Other elements (Calcium, Phosphorus, Potassium, Sulfur, Sodium, Chlorine, Magnesium) are present in smaller amounts.

Trace Elements

• Common in food and water; present in minute quantities in living tissues.

• Important to prevent diseases:

  • Iron is critical for oxygen transport in the body.

  • Iodine deficiency can lead to goiter due to disrupted thyroid hormone production.

Atoms

• Composed of protons, neutrons, and electrons.

• Each element consists of only one kind of atom, being the smallest unit retaining the properties of an element.

Composition of Atoms

• Location:

  • Protons and neutrons are located in the nucleus.

  • Electrons orbit around the nucleus.

• Charge:

  • Protons: positively charged.

  • Electrons: negatively charged.

  • Neutrons: no charge.

Atomic Number and Atomic Mass

• Atomic Number: Number of protons in an atom.

• Mass Number: Sum of protons and neutrons.

• Atomic mass is approximately equal to the mass number.

Isotopes

• Atoms can have the same atomic number but different mass numbers

• Isotopes have the same number of protons, but different neutron counts.

• Radioactive isotopes decay spontaneously, emitting subatomic particles or energy.

Chemical Bonds

Chapter 2.2 Learning Goals

• Explain why atoms share or trade electrons to form bonds.

• Describe differences between ionic, non-polar covalent, polar covalent, and hydrogen bonds.

Electrons and Chemical Properties

• Electrons in outer shells dictate an atom's chemical characteristics.

• There can be one or more electron shells, each at specific distances from the nucleus.

Stability and Bond Formation

• Unfilled electron shells lead to instability.

• The first shell requires 2 electrons; subsequent shells require 8.

Types of Chemical Bonds

• Covalent Bonds (2 types)

  • Non-polar covalent bonds.

  • Polar covalent bonds.

• Ionic Bonds

Covalent Bonds

• Atoms share electrons to form covalent bonds (e.g., H2 molecule).

• Variants include single, double, and triple bonds based on the number of shared electron pairs.

• Electronegativity affects electron sharing:

  • Equal sharing results in nonpolar covalent bonds (e.g., H2).

  • Unequal sharing leads to polar covalent bonds (e.g., H2O).

Ionic Bonds

• Formed from the attraction between positively charged cations and negatively charged anions.

• Example: Sodium chloride (NaCl).

Hydrogen Bonds

• Weak bonds that play a critical role in biological chemistry.

• Form between water molecules due to partial charges.

Properties of Water

Chapter 2.3 Learning Goals

• Explain how water dissolves ionic and polar molecules via hydrogen bonds.

• Describe acids and bases in terms of pH and H+ ion concentration.

Water's Polarity

• Water has partial charges, making it polar.

• Polar substances form favorable hydrogen bonds with water (hydrophilic).

• Nonpolar substances do not interact well with water (hydrophobic).

Water as a Solvent

• Solutions consist of uniform mixtures of substances, with water as the universal solvent (aqueous solutions).

• Soluble interactions exemplified by dissolving sugar in water.

Acid-Base Chemistry

• In water, some molecules break into ions (H+ and OH-).

• Acids: Increase H+ concentration.

• Bases: Decrease H+ concentration.

• pH Scale: Ranges from 0 (most acidic) to 14 (most basic).

  • Acidic (0-6.9), Neutral (7), Basic (7.1-14).

Importance of pH Homeostasis

• Most cells require a pH of 7.2 - 7.6 for optimal function.

• Deviations can disrupt cellular function and may lead to severe health consequences.

Buffers

• Buffers stabilize pH by resisting changes, either by releasing or binding H+ ions.

• Example: Carbonic acid-bicarbonate system is crucial for blood pH regulation.

Summary

• Matter is composed of diverse atomic elements.

• Chemical bonds, including covalent and ionic, are formed through electron interactions.

• Water's polarity facilitates hydrogen bond formation and makes it an exceptional solvent.

• pH serves as a crucial measure of hydrogen ion concentrations, with buffers playing a key regulatory role in biological systems.