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  • Solomon/Berg/Martin Biology, 8e
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Atoms and Molecules: The Chemical Basis of Life

  • Chapter 2

Learning Objective 1

  • What principal chemical elements are found in living things?
  • What are the most important functions of these elements?


  • A substance that cannot be decomposed into simpler substances by normal chemical reactions

The Periodic Table


  • Carbon, hydrogen, oxygen, and nitrogen are the most abundant elements in living things (about 96% of mass)

Element Functions

  • Carbon
    • backbone of organic molecules
  • Hydrogen and oxygen
    • components of water
  • Nitrogen
    • component of proteins and nucleic acids

Learning Objective 2

  • Compare the physical properties (mass and charge) and locations of electrons, protons, and neutrons
  • What is the difference between the atomic number and the mass number of an atom?


  • Nucleus
    • protons (positive)
    • neutrons (uncharged)
  • Electrons
    • encircle the nucleus
    • negative

Atomic Number and Mass

  • Each atom is a particular element
    • identified by number of protons (atomic number)
  • Atomic mass
    • sum of protons and neutrons


  • Atomic Mass Unit
    • Mass of a single proton or neutron
  • Mass of an electron
    • about 1/1800 amu


  • Fig. 2-2, p. 28
  • Carbon-12 (12 C)
  • (6p, 6n)
  • 6
  • Carbon-14 (14C)
  • (6p, 8n)
  • 6

Learning Objective 3

  • What are orbitals and electron shells?
  • How are electron shells related to principal energy levels?


  • Chemical properties of an atom are determined by its highest-energy (valence) electrons


  • Electrons move rapidly in electron orbitals
    • Outside the nucleus
  • Electron shell
    • Electrons in orbitals at the same principal energy level
    • Electron in shell far from nucleus has more energy than electron in shell close to nucleus

Atomic Orbitals

  • Fig. 2-4a, p. 30
  • (a) The first principal energy level contains a maximum of 2 electrons, occupying a single spherical orbital (designated 1s). The electrons depicted in the diagram could be present anywhere in the blue area.
  • Nucleus
  • Fig. 2-4b, p. 30
  • (b) The second principal energy level includes four orbitals, each with a maximum of 2 electrons: one spherical (2s) and three dumbbell-shaped (2p) orbitals at right angles to one another.
  • Fig. 2-4c, p. 30
  • (c) Orbitals of the first and second principal energy levels of a neon
  • atom are shown superimposed. Note that the single 2s orbital plus
  • three 2p orbitals make up neon's full valence shell of 8 electrons.
  • Compare this more realistic view of the atomic orbitals with the
  • Bohr model of a neon atom at right.

Learning Objective 4

  • How does the number of valence electrons of an atom relate to its chemical properties?

Valence Electrons

  • Electron in the outer shell
    • most energetic electrons
  • Number and arrangement of an atom’s valence electrons
    • determine its chemical properties

Valence Electrons

  • An atom tends to lose, gain, or share electrons to fill its valence shell
  • Electrons needed to fill valence shell
    • Most atoms: 8 electrons
    • Hydrogen or helium: 2 electrons

Learning Objective 5

  • What is the difference between simplest, molecular, and structural chemical formulas?


  • Atoms are joined by chemical bonds to form compounds
  • A chemical formula gives the types and relative numbers of atoms in a substance

Chemical Formulae

  • Simplest formula
  • Molecular formula
    • actual numbers of each type of atom
  • Structural formula
    • the arrangement of atoms in a molecule


  • A molecule consists of atoms joined by covalent bonds
  • Other important chemical bonds include ionic bonds and hydrogen bonds

Learning Objective 6

  • Why is the mole concept so useful to chemists?

Avogadro’s Number

  • Avogadro’s number 6.02 x 1023
  • One mole (atomic or molecular mass in grams) of any substance contains 6.02 x 1023 atoms, molecules, or ions
  • Enables scientists to “count” particles by weighing a sample

Learning Objective 7

  • What is the difference between covalent bonds, ionic bonds, hydrogen bonds, and van der Waals interactions?
  • How does each differ in the mechanisms by which they form and in relative strength?

Covalent Bonds

  • Strong, stable bonds
  • Formed when atoms share valence electrons
  • Form molecules
  • May rearrange the orbitals of valence electrons (orbital hybridization)

Covalent Bonds

  • Fig. 2-5, p. 32
  • Molecular hydrogen (H2)
  • H
  • H
  • H
  • H
  • or
  • Hydrogen (H)
  • Hydrogen (H)
  • (a) Single covalent bond formation. Two hydrogen atoms achieve stability by sharing a pair of electrons,
  • thereby forming a molecule of hydrogen. In the structural formula on the right, the straight line between the hydrogen atoms represents a single covalent bond.
  • Molecular oxygen (O2)
  • (double bond is formed)
  • O
  • O
  • O
  • O
  • Oxygen (O)
  • Oxygen (O)
  • or
  • (b) Double covalent bond formation. In molecular oxygen, two oxygen atoms share two pairs of
  • electrons, forming a double covalent bond. The parallel straight lines in the structural formula represent
  • a double covalent bond.

Nonpolar and Polar Covalent Bonds

  • Covalent bonds are
    • nonpolar if electrons are shared equally between the two atoms
    • polar if one atom is more electronegative (greater electron affinity) than the other

Ionic Bonds

Ionic Bonds

  • Fig. 2-9a, p. 35
  • +
  • 11 protons
  • 11 electrons
  • Sodium (Na)
  • 17 electrons
  • Chlorine (Cl)
  • 17 protons
  • and
  • 18 electrons
  • Chloride ion (Cl–)
  • 10 electrons
  • Sodium ion (Na+)
  • Sodium chloride (NaCl)

Hydrogen Bonds

  • Relatively weak bonds
  • Form when
    • A hydrogen atom with a partial positive charge
    • Is attracted to an atom (usually O or N) with a partial negative charge
    • Already bonded to another molecule or part of the same molecule

Hydrogen Bonds

  • Fig. 2-11, p. 35
  • Electronegative
  • atoms
  • Hydrogen bond
  • H
  • H
  • H
  • H
  • H
  • +
  • O
  • N


van der Waals interactions

  • Weak forces
  • Based on fluctuating electric charges

Learning Objective 8

  • What are oxidation and reduction reactions?
  • How do oxidation and reduction reactions relate to the transfer of energy?

Redox Reactions

  • Oxidation and reduction reactions
  • Electrons (energy) are transferred from a reducing agent to an oxidizing agent

Oxidation and Reduction

  • Oxidation
    • Atom, ion, or molecule loses electrons (energy)
  • Reduction
    • Atom, ion, or molecule gains electrons (energy)


  • The energy of an electron is transferred in a redox reaction

Learning Objective 9

  • How do hydrogen bonds between adjacent water molecules govern the properties of water?

Polar Molecules

  • Water is a polar molecule
  • One end has a partial positive charge and the other has a partial negative charge
  • Because it is polar, water is an excellent solvent for ionic or polar solutes

Polar Molecules

  • Fig. 2-7, p. 34
  • Hydrogen parts
  • Hydrogen (H)
  • Oxygen (O)
  • Hydrogen (H)
  • Partial
  • negative
  • charge at
  • oxygen end
  • of molecule
  • Oxygen part
  • Water molecule (H2O)
  • Partial
  • positive
  • charge
  • at hydrogen
  • end of
  • molecule
  • Fig. 2-7, p. 34
  • Hydrogen (H)
  • Oxygen (O)
  • Hydrogen (H)
  • Partial
  • negative
  • charge at
  • oxygen end
  • of molecule
  • Water molecule (H2O)
  • Partial
  • positive
  • charge
  • at hydrogen
  • end of
  • molecule
  • Oxygen part
  • Hydrogen parts
  • Stepped Art

Cohesion and Adhesion

  • Water molecules exhibit cohesion because they form hydrogen bonds with one another
  • Water molecules exhibit adhesion by hydrogen bonding to substances with ionic or polar regions

Hydrogen Bonds in Water

Specific Heat

  • Water has high specific heat
  • Hydrogen bonds must break to raise water temperature
  • Specific heat of water helps
    • organisms maintain relatively constant internal temperature
    • keep large bodies of water (ocean) at a constant temperature

Heat of Vaporization

  • Water has a high heat of vaporization
  • Hydrogen bonds must break for molecules to enter vapor phase
  • Molecules carry heat, causing evaporative cooling


  • Hydrogen bonds between water molecules make ice less dense than liquid water
  • Because ice floats, the aquatic environment is less extreme

Three Phases of Water


  • Water molecules are polar, with partial positive and negative charges
  • Form hydrogen bonds with one another and other charged substances

Learning Objective 10

  • What is the difference between an acid and a base?
  • What are the properties of acids and bases?

Acids and Bases

  • Acids
    • proton (hydrogen ion, H+ ) donors
    • dissociate in solution to yield H+ and an anion
  • Bases
    • proton acceptors
    • dissociate in solution to yield hydroxide ions (OH-)


  • Acids are hydrogen ion donors
  • Bases are hydrogen ion acceptors
  • The pH scale measures the hydrogen ion concentration of a solution

Learning Objective 11

  • How does the hydrogen ion concentration (moles per liter) of a solution relate to its pH value?
  • How do buffers help minimize changes in pH?


  • The negative log of the hydrogen ion (H+) concentration of a solution
    • (measured in moles per liter)

pH of Solutions

  • Neutral solution
    • equal concentrations of H+ and OH-
    • (10 -7 mol/L), pH 7
  • Acidic solution
    • pH less than 7
  • Basic solution
    • pH greater than 7


  • Buffering system
    • based on a weak acid or a weak base
  • Buffer
    • resists changes in pH of a solution when acids or bases are added

Learning Objective 12

  • What is the composition of a salt?
  • Why are salts are important in organisms?


  • Salt
    • a compound in which the hydrogen atom of an acid is replaced by some other cation
  • Salts provide many mineral ions essential for life functions
  • How Atoms Bond

Spheres of Hydration


The pH Scale


The Shell Model of Electron Distribution


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