2 life and chemistry: small molecules. 2 water and the origin of life’s chemistry atoms: the...

2Life and Chemistry:

Small Molecules

2 Life and Chemistry: Small Molecules

• Water and the Origin of Life’s Chemistry

• Atoms: The Constituents of Matter

• Chemical Bonds: Linking Atoms Together

• Chemical Reactions: Atoms Change Partners

• Water: Structure and Properties

• Acids, Bases, and the pH Scale

• Properties of Molecules

2 Water and the Origin of Life’s Chemistry

• The earliest chemical signatures of life on Earth are about 4 billion years old.

• The presence of water, possibly brought by comets striking the Earth, was critical in making conditions suitable for life.

• Environmental conditions conducive to life evolved during the Hadean period.

Figure 2.1 A Geological Time Scale

2 Atoms: The Constituents of Matter

• All matter is composed of atoms.

• Each atom consists of at least one proton and one electron.

• Atoms have mass. The mass comes mostly from the proton and a neutrally charged body called a neutron.

Figure 2.2 The Helium Atom


• Each element contains only one type of atom.

• Information on elements is arranged in logical order in a table called the periodic table.

• The periodic table arranges elements left to right based on their atomic number, and in columns based on similarities in their properties.

Figure 2.3 The Periodic Table (Part 1)

Figure 2.3 The Periodic Table (Part 2)


• Each element has a unique atomic number which is the number of protons found in an atom of the element.

• The mass number is the number of protons plus the number of neutrons.

• The mass number is used as the weight of the atom, in units called daltons.

• Each element has a unique symbol: H is hydrogen, C is carbon, Na is sodium, and Fe is iron.


• All atoms of an element have the same number of protons, but not necessarily the same number of neutrons.

• Atoms of the same element that have different atomic weights are called isotopes.

Figure 2.4 Isotopes Have Different Numbers of Neutrons


• Some isotopes are radioisotopes, which emit energy as alpha, beta, and gamma radiation from their nuclei.

• Radioactive decay transforms the original atom into another atom, usually of another element.


• The region in which an electron travels is called the electron’s orbital.

• The orbitals constitute a series of electron shells, or energy levels, around the nucleus.

• Two electrons at most can occupy each orbital.


• The first shell is the innermost shell and has just one orbital, called the s orbital.

• The s orbital fills first and its electrons have the lowest energy.

• The second shell is next closest to the nucleus and has one s and three p orbitals.

• The second shell can accommodate eight electrons, two per orbital.

Figure 2.6 Electron Orbitals


• The outermost shell of an atom determines how it reacts with other atoms.

• Generally, if eight electrons are in the outer shell, the atom is stable and does not tend to react.

• Atoms which do not have eight electrons in the outermost shell will share, gain, or lose electrons to arrive at a stable state.


• The tendency of atoms to be stable when they have eight electrons in their outermost shells is called the rule of eight, or the octet rule.

• Hydrogen and phosphorus are exceptions to the octet rule.

2 Chemical Bonds: Linking Atoms Together

• A molecule is two or more atoms bonded together.

• A chemical bond is an attractive force that links two atoms together.


• A covalent bond is formed by sharing of a pair of electrons between two atoms.

• In hydrogen molecules (H2), a pair of electrons share a common orbital and spend equal amounts of time around each of the two nuclei.

• The nuclei stay some distance from each other due to mutually repelling positive charges.

Figure 2.8 Electrons Are Shared in Covalent Bonds

Properties of Molecules


• Molecules made up of more than one type of atoms are called compounds.

• Every compound has a molecular weight that is the sum of all atoms in the molecule.


• Covalent bonds are very strong.

• Each covalent bond has a predictable length, angle, and direction, which makes it possible to predict the three-dimensional structures of molecules.

• A double covalent bond occurs when atoms share two pairs of electrons; in triple covalent bonds atoms share three electron pairs.

Figure 2.10 Covalent Bonding With Carbon



• Electrons are not always shared equally between covalently bonded atoms.

• The attractive force that an atom exerts on electrons is called electronegativity.

• When a molecule has nuclei with different electronegativities, an electron spends most of its time around the nucleus with the greater electronegativity.


• Unequal sharing of electrons causes a partial negative charge around the more electronegative atom, and a partial positive charge around the less electronegative atom, resulting in a polar covalent bond.

• Molecules that have polar covalent bonds are called polar molecules.

Figure 2.11 The Polar Covalent Bond in the Water Molecule



• Hydrogen bonds may form within or between atoms with polar covalent bonds.

• The – portion of one molecule has a weak attraction to the + portion of another molecule. Each of these attractions is called a hydrogen bond.

• Hydrogen bonds do not share electrons.

• Although hydrogen bonds are weak, they tend to be additive, and they are of profound biological importance.

Figure 2.12 Hydrogen Bonds Can Form between or within Molecules



• Ionic bonds involve a complete transfer of one or more electrons.

• Ions are formed when an atom loses or gains electrons.

• Positively charged ions are called cations.

• Negatively charged ions are called anions.

Figure 2.13 Formation of Sodium and Chloride Ions



• Ionic bonds are formed by the electrical attraction between ions with opposite charges.

• Table salt has chloride and sodium ions, held together by ionic bonds.

• When salt is introduced into water, the partial charges of the water molecules can easily interfere with the ionic bonds.

Figure 2.14 Water Molecules Surround Ions



• Polar molecules tend to be hydrophilic. Substances that are ionic or polar often dissolve in water due to hydrogen bonds.

• Nonpolar molecules are called hydrophobic because they tend to aggregate with other nonpolar molecules.

• Nonpolar molecules are also attracted to each other via relatively weak attractions called van der Waals forces.

2 Chemical Reactions: Atoms Change Partners

• Chemical reactions occur when atoms combine or change partners.

• In a chemical reaction, reactants are converted to products.

• A chemical reaction can be written as an equation. The equation must balance because matter is neither created nor destroyed.

Figure 2.15 Bonding Partners and Energy May Change in a Chemical Reaction


2 Chemical Reactions: Atoms Change Partners

• Changes in energy usually accompany chemical reactions.

• Stored energy, such as that in chemical bonds, is called potential energy and is available for future use.

• We can measure the potential energy of molecules and express it in units of heat called calories.

• A calorie is the amount of heat required to raise the temperature of one gram of pure water from 14.5°C to 15.5°C.

2 Water: Structure and Properties

• Due to its shape, polarity, and ability to form hydrogen bonds, water has some unusual properties.


• Ice is held in a crystalline structure by the orientation of water molecules’ hydrogen bonds.

• Each molecule forms hydrogen bonds with four other molecules.

• These four hydrogen bonds increase the space the water molecules take up, so water expands as it freezes, and ice is less dense than liquid water.

• For these reasons, ice floats in liquid water.

Figure 2.16 Hydrogen Bonds Hold Water Molecules Together (Part 1)


Figure 2.16 Hydrogen Bonds Hold Water Molecules Together (Part 2)



• Compared to other nonmetallic substances, ice requires a lot of heat to melt because hydrogen bonds must be broken.

• The opposite process, freezing, requires water to lose a great deal of heat.


• A great deal of heat energy is required to change the temperature of liquid water because the hydrogen bonds must be broken.

• Specific heat is the number of calories needed to raise one gram of a substance 1oC. The specific heat of liquid water is 1.

• Liquid water has a higher specific heat than most other small molecules in liquid form.


• The heat of vaporization is the amount of heat needed to change a substance from its liquid state to its gaseous state.

• A lot of heat is required to change water to a gaseous state because the hydrogen bonds of the liquid water must be broken.

• Evaporation has a cooling effect by absorbing calories.

• Condensing has the opposite effect, releasing heat.


• Water has a cohesive strength because of hydrogen bonds.

• The cohesive strength of water molecules allows the transport of water from the roots to the tops of trees.

• Water has high surface tension, which means that the surface of liquid water is relatively difficult to puncture.


• Water is the solvent of life.

• Living organisms are over 70 percent water by weight and many reactions take place in this watery environment.

• A solution is a substance (the solute) dissolved in a liquid (the solvent).


• Reactions that take place in aqueous solutions may be studied in two ways:

Qualitative analysis is the study of substances dissolved in a solvent and their reactions.

Quantitative analysis measures the amounts of substances and solvents.


• The mole concept is fundamental to quantitative analysis. A mole is the amount of a substance in grams whose weight is equal to its molecular weight.

• One mole of any given compound contains approximately 6.03 x 1023 molecules of that compound (Avogadro’s number).


• A 1 molar (1 M) solution is one mole of a compound dissolved in water to make one liter.

• Example: One mole of NaCl is the atomic weight of Na (23) plus the atomic weight of Cl (35.5), or 58.5, in grams. When 58.5 grams of NaCl are dissolved in water to make one liter, the solution is 1 molar.

2 Acids, Bases, and the pH Scale

• Some substances dissolve in water and release hydrogen ions (H+); these are called acids. Their release is called ionization.

• Other substances dissolve in water and release hydroxide ions (OH–); these are called bases.

• Acids donate H+; bases accept H+.


• Acids release H+ ions in solution.

• If the reaction is complete, it is a strong acid, such as HCl.

• The carboxyl group (—COOH) is common in biological compounds. It functions as an acid because

—COOH —COO– + H+


• Bases accept H+ in solution.

• NaOH ionizes completely to Na+ and OH–. The OH– absorbs H+ to form water. It is a strong base.

• The amino group (—NH2) is an important part of many biological compounds; it functions as a weak base by accepting H+:

—NH2 + H+ —(NH3)+


• Ionization of strong acids is virtually irreversible.

• Ionization of weak acids and bases is somewhat reversible.

• Many large molecules in biological systems contain weak acid or base groups.


• Water is really a weak acid and has a slight tendency to ionize into H+ and OH–.

• This ionization is very important for living creatures and the chemical reactions they must perform because the H+ ion is so reactive.


• pH is the measure of hydrogen ion concentration.

• It is defined as the negative logarithm of the hydrogen ion concentration in moles per liter.

• The pH scale indicates the strength of a solution of an acid or base. The scale values range from 1 through 14.

• A pH 7 means the concentration of hydrogen ions is 1 x 10–7 moles per liter of water.

Figure 2.18 pH Values of Some Familiar Substances



• A buffer is a mixture of a weak acid and its corresponding base.

• Because buffers can react with both added bases and acids, they make the overall solution resistant to pH change.

• Buffers illustrate the law of mass action: Addition of reactants to one side of a reaction

drives the reaction in the direction that uses that component.

Figure 2.19 Buffers Minimize Changes in pH


2 Properties of Molecules

• Chemists use the characteristics of composition, structure, reactivity, and solubility to help classify molecules.

• Two other properties that influence the behavior of molecules are the presence of recognizable functional groups, and the existence of isomers of molecules.


• Functional groups give specific properties to molecules.

• Functional groups are covalently bonded to organic molecules.

• Amino acids are biological molecules that contain both a carboxyl group and an amino group.

Figure 2.20 Some Functional Groups Important to Living Systems (Part 1)



• Isomers are molecules that have the same chemical formula but different arrangements of the atoms.

• Structural isomers differ in terms of how atoms are joined together.

• Optical isomers are mirror images of each other.

• Optical isomers can occur whenever a carbon has four different atoms or groups attached to it.

Figure 2.21 Optical Isomers


2 life and chemistry: small molecules. 2 water and the origin of life’s chemistry atoms: the...

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