Chapter 2: The Chemical Context of Life – Study Notes

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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 takes up space and has mass. Matter is made up of elements, and elements can combine to form compounds with unique properties.

Matter: Anything that occupies space and has mass.
Element: A substance that cannot be broken down to other substances by chemical reactions (e.g., oxygen, carbon).
Compound: A substance consisting of two or more elements in a fixed ratio (e.g., water, H2O; sodium chloride, NaCl).
Emergent Properties: Compounds have characteristics different from those of their constituent elements.

Example: Sodium (Na, a reactive metal) and chlorine (Cl, a poisonous gas) combine to form sodium chloride (NaCl, table salt), which is safe to eat.

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, such as iron (Fe), zinc (Zn), and iodine (I).

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

Each element consists of unique atoms, which are the smallest units of matter that retain the properties of the element. Atoms are composed of subatomic particles: neutrons, protons, and electrons.

Neutrons: No electrical charge; found in the nucleus.
Protons: Positive charge; found in the nucleus.
Electrons: Negative charge; orbit the nucleus in electron shells.
Atomic Nucleus: Center of the atom, containing protons and neutrons.
Dalton: Unit of mass for atoms and subatomic particles (1 dalton ≈ mass of 1 proton or neutron).

Atomic Number and Atomic Mass

Atomic Number (Z): Number of protons in the nucleus; defines the element.
Mass Number (A): Sum of protons and neutrons in the nucleus.
Atomic Mass: Total mass of an atom, approximately equal to the mass number.

Example: Nitrogen has 7 protons (atomic number 7) and typically 7 neutrons (mass number 14).

Isotopes

Isotopes: Atoms of the same element with different numbers of neutrons.
Radioactive Isotopes: Unstable isotopes that decay spontaneously, emitting particles and energy.

Example: Carbon-12, Carbon-13, and Carbon-14 are isotopes of carbon, differing in neutron number.

Table: Subatomic Particles in Selected Elements

Element	Protons (P)	Neutrons (N)	Electrons (E)	Atomic # (A.#)	Atomic Mass (A.M.)
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
Calcium2+	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 arranges elements by their electron configurations.

Valence Electrons: Electrons in the outermost shell (valence shell); determine chemical reactivity.
Chemically Inert: Elements with a full valence shell (e.g., noble gases) are unreactive.

Example: Helium (He) has a full valence shell and is chemically 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 valence electrons, resulting in chemical bonds that hold atoms together in molecules or compounds.

Covalent Bonds

Covalent Bond: Sharing of a pair of valence electrons between two atoms.
Single Bond: Sharing of one pair of electrons (e.g., H–H).
Double Bond: Sharing of two pairs of electrons (e.g., O=O).
Molecule: Two or more atoms held together by covalent bonds.
Structural Formula: Shows arrangement of atoms and bonds (e.g., H–H).
Molecular Formula: Indicates the number and type of atoms (e.g., H2).

Example: Water (H2O) is a molecule formed by covalent bonds between hydrogen and oxygen.

Electronegativity and Bond Polarity

Electronegativity: An atom’s attraction for electrons in a covalent bond.
Nonpolar Covalent Bond: Electrons are shared equally (e.g., H2).
Polar Covalent Bond: Electrons are shared unequally, creating partial charges (e.g., H2O).

Example: In water, oxygen is more electronegative than hydrogen, resulting in a polar molecule with partial charges.

Ionic Bonds

Atoms may transfer electrons, resulting in charged particles called ions.
Cation: Positively charged ion (e.g., Na+).
Anion: Negatively charged ion (e.g., Cl–).
Ionic Bond: Attraction between a cation and an anion.
Ionic Compounds (Salts): Compounds formed by ionic bonds, often found as crystals (e.g., NaCl).

Example: Sodium (Na) transfers an electron to chlorine (Cl), forming NaCl (table salt).

Hydrogen Bonds

Hydrogen Bond: Weak attraction between a hydrogen atom covalently bonded to one electronegative atom and another electronegative atom.
Important in the structure of water, proteins, and DNA.

Example: Hydrogen bonds form between water molecules, contributing to water’s unique properties.

Additional info: Understanding atomic structure and chemical bonding is foundational for studying biological molecules and processes.