Chapter 2: The Chemical Context of Life – Study Notes

Chapter 2: The Chemical Context of Life

Concept 2.1: Matter Consists of Chemical Elements in Pure Form and in Combinations Called Compounds

All living organisms are composed of matter, which is anything that occupies space and has mass. Matter is made up of elements and compounds, each with distinct properties and roles in biological systems.

• Matter: Anything that takes up space and has mass.

• Element: A substance that cannot be broken down to other substances by chemical reactions.

• Compound: A substance consisting of two or more elements in a fixed ratio. Compounds have characteristics different from those of their constituent elements.

• Example: Sodium (Na) and chlorine (Cl) combine to form sodium chloride (NaCl), which has properties distinct from either element.

Table: Elements in the Human Body

Element	Symbol	Percentage of Body Mass (including water)
Oxygen	O	65.0%
Carbon	C	18.5%
Hydrogen	H	9.5%
Nitrogen	N	3.3%
Calcium	Ca	1.5%
Phosphorus	P	1.0%
Potassium	K	0.4%
Sulfur	S	0.3%
Sodium	Na	0.2%
Chlorine	Cl	0.2%
Magnesium	Mg	0.1%

Additional info: Trace elements (less than 0.01% of mass) are also essential for life, including boron, chromium, copper, fluorine, iodine, iron, manganese, molybdenum, selenium, silicon, tin, vanadium, and zinc.

Concept 2.2: An Element's Properties Depend on the Structure of Its Atoms

Each element is composed of unique atoms, which are the smallest units of matter that retain the properties of the element. The structure of atoms determines their chemical behavior.

• Atom: The smallest unit of matter that retains the properties of an element.

• Subatomic Particles: Atoms are made of neutrons (no charge), protons (positive charge), and electrons (negative charge).

• Atomic Nucleus: Neutrons and protons form the nucleus; electrons form a cloud around the nucleus.

• Dalton: Unit of measurement for atomic and subatomic particles.

Atomic Number and Atomic Mass

• Atomic Number: Number of protons in the nucleus.

• Mass Number: Sum of protons and neutrons in the nucleus.

• Atomic Mass: Total mass of the atom, approximated by the mass number.

Example: Nitrogen has atomic number 7 and mass number 14.

Isotopes

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Radioactive Isotopes: Decay spontaneously, emitting particles and energy.

• Example: Carbon-12, Carbon-13, and Carbon-14 are isotopes of carbon.

Table: Subatomic Composition of Selected Elements

Element	P (Protons)	N (Neutrons)	E (Electrons)	A.# (Atomic Number)	A.M. (Atomic Mass)
Hydrogen	1	0	1	1	1
Helium	2	2	2	2	4
Carbon	6	6	6	6	12
Carbon-14	6	8	6	6	14
Nitrogen	7	7	7	7	14
Calcium 2+	20	20	18	20	40

Electron Distribution and Chemical Properties

The chemical behavior of an atom is determined by the distribution of electrons in electron shells. The periodic table organizes elements by their electron configurations.

• Valence Electrons: Electrons in the outermost shell (valence shell).

• Chemically Inert: Elements with a full valence shell do not readily react with other elements.

• Example: Helium has a full valence shell and is inert.

Concept 2.3: The Formation and Function of Molecules Depend on Chemical Bonding Between Atoms

Atoms with incomplete valence shells can share or transfer electrons, resulting in chemical bonds that hold atoms together in molecules and compounds.

• Chemical Bond: Attraction that holds atoms together.

Covalent Bonds

• Covalent Bond: Sharing of a pair of valence electrons between two atoms.

• Single Bond: Sharing of one pair of electrons.

• Double Bond: Sharing of two pairs of electrons.

• Molecule: Two or more atoms held together by covalent bonds.

• Structural Formula: Representation of atoms and bonds (e.g., H—H).

• Molecular Formula: Abbreviated representation (e.g., H2).

Electronegativity and Bond Polarity

• Electronegativity: An atom’s attraction for electrons in a covalent bond.

• Nonpolar Covalent Bond: Electrons are shared equally.

• Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges.

• Example: Water (H2O) is a polar molecule.

Ionic Bonds

• Ionic Bond: Attraction between oppositely charged ions (cation and anion).

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Example: Sodium chloride (NaCl) is formed by the transfer of an electron from sodium to chlorine.

Table: Ionic Bond Formation

Atom	Electron Transfer	Ion Formed
Na (Sodium)	Loses 1 electron	Na+ (Cation)
Cl (Chlorine)	Gains 1 electron	Cl- (Anion)

Hydrogen Bonds

• Hydrogen Bond: Weak attraction between a hydrogen atom covalently bonded to one electronegative atom and another electronegative atom.

• Example: Hydrogen bonds occur between water molecules and between water and ammonia.

Additional info: Hydrogen bonds are crucial for the structure and function of biological molecules, such as DNA and proteins.