Chapter 2 – The Chemical Basis of Life I: Atoms, Molecules, and Water
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Chapter 2: The Chemical Basis of Life I: Atoms, Molecules, and Water
Chapter Outline: 1) Atoms 2) Chemical bonds and molecules 3) Chemical reactions 4) Properties of water 5) pH and buffers
Note: Blue font color marks important ideas; bold marks terms to master.
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2.1 Atoms Learning Outcomes
Sketch the general structure of atoms.
Define orbital and electron shell.
Relate atomic structure to the periodic table of the elements.
Quantify atomic mass using units of daltons and moles.
Explain how a single element may exist in 2 or more forms, called isotopes.
List the 4 elements that make up most of the mass of all living organisms (O, C, H, N).
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2.1 Atoms
Atoms are the smallest functional units of matter that form all chemical substances.
Atoms cannot be further broken down into other substances by ordinary chemical or physical means.
An element is a pure substance made up of only 1 kind of atom.
When 2 or more atoms are bonded together, a molecule is formed.
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2.1 Atoms — Subatomic Particles
Three subatomic particles are found within an atom:
Protons (positive charge, +) — located in the atomic nucleus.
Neutrons (neutral) — located in the atomic nucleus.
Electrons (negative charge, −) — located in orbitals around the nucleus.
Protons and electrons are typically present in equal numbers, giving the atom no net charge.
The number of neutrons can vary for a given atom (isotopes differ in neutron number).
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2.1 Atoms — Orbitals Around the Nucleus
Electrons occupy three-dimensional volumes of space called orbitals that surround the nucleus.
Each orbital can hold a maximum of 2 electrons.
Orbitals are organized into electron shells; shells have characteristic energy levels.
Shells are numbered from the nucleus outward (shell 1 is closest).
A shell can contain 1 or more orbitals.
Example: the first shell contains 1 orbital (max 2 electrons);
the second shell contains 4 orbitals (max 8 electrons).
Note: 2.3 figure reference: Fig 2.3, Biology, Brooker.
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2.1 Atoms — Filling Orbitals and Valence Electrons
Electrons fill orbitals in a specific order, starting with the shell closest to the nucleus.
Valence electrons are electrons in the outermost shell and can participate in chemical bonds.
Note: 2.4 figure reference: Fig 2.4, Biology, Brooker.
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2.1 Atoms — Atomic Number
Each element has a unique number of protons, called its atomic number (Z).
The atomic number distinguishes one element from another and equals the number of electrons in a neutral atom, giving no net charge.
The periodic table is arranged according to atomic number and electron shells.
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2.1 Atoms — Periodic Table Organization
Rows (periods) correspond to the number of electron shells.
Columns (groups) indicate the number of electrons in the outer shell (valence electrons).
Elements in the same column have similar properties due to the same number of valence electrons.
Figure reference: Fig 2.5, Biology, Brooker.
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2.1 Atoms — Atomic Mass and Daltons
Protons and neutrons have masses close to each other, both > 1,800 times the mass of an electron; electron mass is negligible.
Atomic mass is measured in daltons (Da); 1 Da = 1/12 the mass of a carbon-12 atom.
Some sources use atomic mass units (amu); 1 Da = 1 amu.
Practically: proton ≈ 1 Da, neutron ≈ 1 Da, electron mass ≈ 0 Da.
Practical question: atomic mass of an atom with 4 protons, 5 neutrons, and 4 electrons is about 4 + 5 = 9 Da (electrons contribute negligibly).
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2.1 Atoms — The Mole
Some atoms have vastly different masses; 1 gram of element can contain many or few atoms.
The mole measures the number of particles: { mole}=6.022x10^{23}{ particles}.
The mass in grams of 1 mole of a substance equals its atomic or molecular mass (in Da).
2.1 Atoms — Isotopes
Isotopes differ in neutron number but have the same atomic number.
Example: Carbon-12 (12C) has 6 protons and 6 neutrons; Carbon-14 (14C) has 6 protons and 8 neutrons.
Atomic masses on the periodic table are average masses of naturally occurring isotopes.
Some isotopes are unstable radioisotopes that emit energy/radiation as they decay.
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2.1 Atoms — Major Elements
Living organisms are largely composed of four elements: Oxygen (O), Carbon (C), Hydrogen (H), and Nitrogen (N).
Minerals and trace elements are also required for growth and normal function.
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2.2 Chemical Bonds and Molecules — Learning Outcomes
Compare and contrast bond types and atomic interactions: covalent, ionic, hydrogen, van der Waals.
Explain electronegativity and its role in polar vs nonpolar covalent bonds.
Describe how a molecule’s shape influences interactions with other molecules.
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2.2 Molecules and Compounds
A molecule contains 2 or more atoms bonded together.
A compound is a molecule that contains different kinds of atoms.
Molecular formulas represent molecules; chemical symbols show elements, with subscripts indicating atom counts (e.g. C6H12O6).
Emergent properties: compounds often have properties different from their constituent elements.
Bonds hold atoms together: covalent (polar & nonpolar), hydrogen bonds, ionic bonds are key in biology.
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2.2 Covalent Bonds — Sharing Electrons
Covalent bonds form when atoms share a pair of electrons to fill outer shells.
Occur between atoms with unfilled valence shells; stability increases when the valence shell is full.
The octet rule: many atoms aim to have 8 electrons in the outer shell; hydrogen is an exception (fills with 2 electrons).
Covalent bonds are strong and stable.
Bond types by electron sharing:
1 pair of electrons — single bond
2 pairs of electrons — double bond
3 pairs of electrons — triple bond
Examples follow for different bonding scenarios.
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2.2 Covalent Bonds — Carbon and Bonding Capacity
Each atom forms a characteristic number of covalent bonds based on outer-shell needs.
Carbon can form 4 covalent bonds, enabling it to link to multiple atoms and form the backbone of biological macromolecules (carbohydrates, lipids, proteins, nucleic acids).
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2.2 Electronegativity — Bond Polarity
Electronegativity is a measure of an atom’s ability to attract shared electrons.
When electronegativities differ, bonds can be polar or nonpolar.
Example electronegativity values (approximate):
H = 2.20
C = 2.55
N = 3.04
O = 3.44
Na = 0.93
Cl = 3.16
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2.2 Polar vs Nonpolar Covalent Bonds
Nonpolar covalent bonds: atoms with similar electronegativities share electrons equally; no charge difference across the molecule.
Polar covalent bonds: atoms with different electronegativities share electrons unequally; results in partial charges (polarity).
Example: In water, O–H bonds are polar covalent due to higher electronegativity of O, creating partial charges: δ- on O and δ+ on H.
Consequence: polar bonds create dipoles and affect molecule interactions.
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2.2 Polar and Nonpolar Molecules
Nonpolar molecules: predominantly nonpolar bonds (e.g., {C–C},{C–H} ).
Polar molecules: contain many polar bonds (e.g., {O–H},{N–H},{O–C} ).
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2.2 Hydrogen Bonds and van der Waals Forces
Hydrogen bonds: occur when a hydrogen atom with a partial positive charge (δ+) interacts with an electronegative atom/molecule (e.g., O, N).
Represented by dashed lines; covalent bonds are solid lines.
Individual hydrogen bonds are weak, but many can sum to strong interactions in aggregate.
Important in protein and DNA structure.
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2.2 van der Waals Dispersion Forces
Weaker than hydrogen bonds; arise in nonpolar molecules.
Caused by temporary, uneven electron distribution leading to brief attractions.
Collective van der Waals forces can contribute significantly to molecular interactions (e.g., in dense packing).
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2.2 Ionic Bonds
Ions: atoms or molecules that have gained or lost electrons to achieve a full valence shell.
Cations: positively charged (+).
Anions: negatively charged (−).
Ionic bond: electrostatic attraction between a cation and an anion.
Ionic compounds are called salts (e.g., NaCl,KCl,CaCl_2).
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2.2 Molecular Shapes and Flexibility
Molecules can adopt different shapes without breaking covalent bonds.
Rotation around bonds can lead to different conformations and shapes.
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2.3 Chemical Reactions Learning Outcomes
Define a chemical reaction and provide an example.
Relate chemical reactions to chemical equilibrium.
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2.3 Chemical Reactions — Fundamentals
A chemical reaction occurs when one or more substances are changed into other substances by making/breaking chemical bonds.
Properties of chemical reactions:
Require energy input (e.g., heat) to mobilize molecules for encounters.
Often catalyzed by enzymes in living organisms to speed up reaction rates.
Tend to proceed in a particular direction but will eventually reach equilibrium.
In living organisms, many reactions occur in watery environments.
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2.3 Chemical Reactions — Reactants, Products, and Equilibrium
Starting materials: reactants; end materials: products.
Example reaction: ext{CH}4 + 2 ext{O}2
ightleftharpoons ext{CO}2 + 2 ext{H}2 ext{O}.The bidirectional arrow indicates the reaction can proceed in both directions.
Equilibrium: rate of product formation equals rate of reactant formation.
In biological systems, most reactions do not reach true equilibrium because products are often quickly consumed in subsequent reactions.
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2.4 Properties of Water — Learning Outcomes
Define solute and solvent.
Compare hydrophilic and hydrophobic substances.
Explain how solution concentration is quantified using molarity.
Describe the 3 states of H2O.
List and explain roles of water critical for life.
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2.4 Properties of Water — Prevalence
Water is a major component of organisms.
Up to 95% of some plants’ weight can be from water;
Typical human body weight is 60–70% water.
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2.4 Water — Solubility and Solutions
In a solution, solute is dissolved in solvent; the solvent is the dissolving medium.
In aqueous solutions, water is the solvent.
Hydrophilic substances readily dissolve in water due to interactions with water’s partial charges.
Hydrophilic substances interact with water via electrical attractions.
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2.4 Hydrophilic vs Hydrophobic, Amphipathic
Hydrophobic substances do not dissolve in water; typically nonpolar (carbon and hydrogen rich).
Oil is a classic hydrophobic substance.
Amphipathic molecules possess both polar/ionized and nonpolar regions, enabling unique interactions with water.
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2.4 Amphipathic Molecules and Micelles/Bilayers
Amphipathic molecules can form micelles or bilayers when in water.
Nonpolar (hydrophobic) regions orient toward the center; polar (hydrophilic) regions face the surface.
These arrangements promote stable chemical interactions and interactions with other polar/nonpolar structures.
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2.4 Concentration in Solutions
Solute concentration: amount of solute per unit volume of solution.
Units can be grams per liter (g/L) or moles per liter (M).
Example: dissolving 1 g NaCl in enough water to make 1 liter yields a concentration of 1 ext{ g/L}.
However, a 1 g/L glucose solution would have far fewer glucose molecules than 1 g/L NaCl because glucose has greater molecular mass.
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2.4 Molarity and Examples
Molarity (M) defined as the number of moles of solute dissolved in 1 liter of solution
Example comparisons for common substances:
Glucose: molecular mass ≈ 180 Da; 1 M glucose solution contains 6.022 × 10^{23} glucose molecules and requires about 180 g of glucose per liter.
NaCl: molecular mass ≈ 58.4 Da; 1 M NaCl solution would contain 6.022 × 10^{23} NaCl formula units and would require about 58.4 g per liter.
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2.4 States of Water and Physical Properties
Water exists in three states: solid (ice), liquid (water), and gas (water vapor).
State changes involve energy input or release (fusion, vaporization, etc.).
Water is unusually stable as a liquid due to:
High heat of vaporization, high heat of fusion, and high specific heat.
Specific heat: amount of heat required to raise the temperature of 1 g by 1°C.
Hydrogen bonds contribute to these properties.
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2.4 Key Functions of Water in Living Organisms
Solvent for chemical reactions.
Participates in chemical reactions.
Provides mechanical support.
Assists in removing toxic waste components.
Enables evaporative cooling.
Supports cohesion and adhesion, surface tension, and lubrication.
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2.5 pH and Buffers — Learning Outcomes
Explain how H2O dissociates into hydroxide ions (OH−) and hydrogen ions (H+), and calculate ion concentrations at a given pH.
Explain the relationship between hydrogen ion concentration and pH value.
Provide examples of how buffers keep body fluids stable.
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2.5 pH and Buffers — Acid–Base Dissociation
Water spontaneously dissociates to form OH− and H+: H2O → H+ + OH-
In pure water, both [H+] and [OH−] are 10-7 M at 25°C
^{-14} ext{ M}^2.$$
Substances (acids/bases) can release or absorb H+ or OH−, altering these concentrations.
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2.5 pH and Buffers — Acids and Bases
An acid releases H+ when added to a solution; acids increase [H+]; acids are proton donors.
Acids can be strong or weak depending on the degree of dissociation and H+ production.
A base accepts H+ when added to a solution; bases decrease [H+]; some bases release OH− while others bind H+.
Bases are proton acceptors and can be strong or weak depending on dissociation and H+ acceptance.
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2.5 pH and Buffers — pH Scale
pH is a measure of H+ concentration: pH= -log10 [H+]
pH values range from 0 to 14; pH and [H+] are inversely related.
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2.5 pH and Buffers — pH Scale Details
Pure water is neutral with pH = 7 (equal [H+] and [OH−]).
Acidic solutions have pH < 7 (More H+) ; basic solutions have pH > 7 (More OH-) .
The pH scale is logarithmic: moving 1 unit on the scale represents a 10-fold change in [H+].
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2.5 pH and Buffers — Biological Importance of pH
pH can affect molecular shapes, reaction rates, binding interactions, and solubility.
Organisms regulate body fluid pH to stay within a narrow range (e.g., human blood: 7.35–7.45).
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2.5 pH and Buffers — Buffers
A buffer resists pH changes when excess acid or base is added.
Buffers can accept H+ from excess acid or donate H+ to neutralize excess base.
An acid–base buffer system can shift to adjust [H+] in response to pH changes.
Example: the carbonic acid/bicarbonate buffer system is widely used in body fluids.
Buffer reaction (illustrative): CO2 +H2O ←→ H2CO3 ←→ H+ + HCO3 -
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Chapter 2 Summary — 2.1 Atoms; 2.2 Chemical Bonds and Molecules
2.1 Atoms: Atoms are composed of subatomic particles; electrons occupy orbitals around a nucleus; each element has a unique number of protons; atoms have a small but measurable mass; isotopes vary in neutron number; four key elements (O, C, H, N) dominate living matter.
2.2 Chemical bonds and molecules: Covalent bonds form when atoms share electrons to fill outer shells; bonds can be polar or nonpolar due to electronegativity; hydrogen bonds and van der Waals forces enable interactions; ionic bonds involve attractions between positive and negative ions.
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Chapter 2 Summary — 2.3–2.5
2.3 Chemical Reactions: Reactions create new compounds; require energy; often enzyme-catalyzed; tend toward equilibrium; in biology, reactions occur in watery environments.
2.4 Properties of Water: Ions and polar molecules dissolve in water; amphipathic regions; solute concentration quantified by molarity; water exists in three states; water performs numerous life-sustaining tasks.
2.5 pH and Buffers: H+ concentrations are altered by acids/bases; pH measures H+ concentration; buffers minimize pH fluctuations in organisms.