Chapter 2 The Chemistry of Life Chapter Outline

Chapter 2 The Chemistry of Life Chapter Outline Module 2.1 Atoms and Elements (Figures 2.1–2.2) A. Define matter: List the three states in which matt...
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Chapter 2 The Chemistry of Life Chapter Outline

Module 2.1 Atoms and Elements (Figures 2.1–2.2) A. Define matter: List the three states in which matter can exist: , or

,

. Define chemistry:

B. Atoms and Atomic Structure (Figure 2.1): 1. Define atom:

2. Subatomic particles exist in 3 basic forms: protons, neutrons, electrons. a. Protons (p+) are found in the central core of the atom known as the atomic nucleus. What charge do protons carry? b. Neutrons (n0), also found in the atomic nucleus, are slightly larger than protons. What charge do neutrons carry? c. Electrons (e-) are found outside the atomic nucleus. What charge do electrons carry? d. All atoms are electrically neutral, meaning they have no charge. The number of protons and electrons are equal, cancelling out each subatomic particles charge. The number of neutrons does not have to be the same as the protons. 3. Electron shells, regions surrounding the atomic nucleus where the likelihood or probability that an electron may exist, can hold a certain number of electrons. a. The 1st shell closest to the nucleus can hold

electrons.

b. The 2nd shell can hold

electrons.

c. The 3rd shell can hold

electrons but is satisfied with 8.

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d. Some atoms may have more than 3 shells. C. Elements in the Periodic Table and the Human Body (Figure 2.2) 1. The number of

that an atom has in its nucleus is its

atomic number. 2. The atomic number defines an element. a. An element is a substance that cannot be broken down into a simpler substance by chemical means. b. Each element is made of atoms with the same number of protons. 3. The periodic table of the elements lists the elements by their increasing atomic numbers. a. This organizes elements into groups with certain properties. b. Each element is represented by a chemical symbol. 4. What are the four major elements that make up the human body? ,

,

, and

. The human body is also made of 7 mineral elements and 13 trace elements. D. Isotopes and Radioactivity: 1. The mass number is equal to the sum of all the

and

found in the atomic nucleus. 2. An isotope is an atom with the same of number of

, but different

number and same number number and different

.

3. Radioisotopes are unstable isotopes have high energy or radiation that can be released by radioactive decay. This allows the isotope to assume a more stable form. Module 2.2 Matter Combined: Mixtures and Chemical Bonds (Figures 2.3–2.7; Tables 2.1-2.2) A. Matter can be combined physically to form a mixture. What is a mixture?

B. Mixtures (Figure 2.3): List and describe the 3 types of mixtures.

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1.

2.

3.

C. Chemical Bonds: Matter can be combined chemically to form a molecule, which is when the atoms of two or more elements are combined by forming chemical bonds. A chemical bond is not a physical structure, but rather an energy relationship or an attractive force between atoms 1. Define molecule.

Define compound.

2. Very large molecules composed of many atoms are called macromolecules. 3. Molecular formulas are a way to represent molecules symbolically with letters and numbers to show the kinds and numbers of atoms in a molecule. D. Chemical bonds are formed when valence electrons in the valence shell, the outermost electron shell, of atoms interact. 1. Valence electrons determine how an atom interacts with other atoms and whether it will form bonds with a specific atom. 2. Define the octet rule.

3. Define the duet rule.

E. Ions and Ionic Bonds (Figure 2.4): An ionic bond is formed when electrons are from a metal atom to a nonmetal atom and results in the formation of ions: cations and anions. The attraction between opposite charges holds or bonds the atoms to one another, forming a compound called a salt.

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1. What is a cation?

2. What is an anion?

F. Covalent Bonds (Figures 2.5, 2.6; Table 2.1): Covalent bonds, the strongest bond, form when two or more nonmetals

electrons between

themselves. 1. Two atoms can share one (single bond), two (double bond), or three (triple bond) electron pairs. 2. All elements have protons that can attract electrons, a property known as electronegativity. a. An element’s electronegativity increases from the bottom left to the upper right of the periodic table, making fluorine (F) the most electronegative element. b. The more electronegative an element the more strongly it attracts electrons, pulling them away from less electronegative elements. 3. Nonpolar covalent bonds result when two nonmetals in a molecule with similar or identical electronegativities pull with the same force and share the electrons equally. 4. Nonpolar molecules occur in 3 situations: a.

(Figure 2.6a)

b.

c. G. Polar covalent bonds form polar molecules when nonmetals with different electronegativities interact, resulting in an unequal sharing of electrons (Figure 2.6b). 1. The atom with the higher electronegativity becomes partially (d-) as it pulls and holds the shared electrons close to itself. 2. The atom with the lower electronegativity becomes partially (d+) as it allows the shared electrons to be pulled away toward the other atom.

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3. Polar molecules with partially positive and partially negative ends are known as dipoles. H. Hydrogen bonds are

attractions between the partially positive end

of one dipole and the partially negative end of another dipole (Figure 2.7). 1. Hydrogen bonds are responsible for a key property of water: surface tension. 2. Where air and water meet, the polar water molecules are more strongly attracted to one another than they are to nonpolar air molecules. Module 2.3 Chemical Reactions (Figures 2.8–2.10) A. Chemical Notation: Chemical notation is a series of symbols and abbreviations that is used to demonstrate what occurs in a reaction. The chemical equation, the basic form of chemical notation, has two parts: 1.

2.

B. Energy and Chemical Reactions: A chemical reaction has occurred every time a chemical bond is formed, broken, or rearranged, or when electrons are transferred between two or more atoms (or molecules). 1. Define energy.

2. Two general forms of energy are: a.

energy is stored, but can be released to do work at some later time.

b.

energy is potential energy that has been released or set in motion to perform work. All atoms have kinetic energy as they are in constant motion and the faster they move the greater that energy.

3. Energy is found in 3 forms in the human body: 1) , and 3)

, 2) , each of which may be

potential or kinetic depending on the location or process. Describe each type of energy.

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a. Chemical energy

b. Electrical energy

c. Mechanical energy

4. Energy, inherent in all chemical bonds, must be invested any time a chemical reaction occurs. Describe the following reactions: a. Endergonic reactions

b. Exergonic reactions

C. Homeostasis and Types of Chemical Reactions: Three fundamental processes that occur in the body to maintain homeostasis, breaking down molecules, converting the energy in food to a usable form, and building new molecules, are carried out by one of three basic types of chemical reactions: 1. Catabolic reactions (decomposition reactions) occur when a. The general chemical notation for this reaction is AB A+B. b. These are usually exergonic because chemical bonds are broken. 2. Exchange reactions occur when a. The general chemical notation for this reaction is AB + CD AD + BC. b. Oxidation-reduction reactions (redox reactions), a special kind of exchange reaction, occur when electrons and energy are exchanged instead of atoms. The reactant that loses electrons is oxidized while the reactant that gains electrons is reduced. c. Redox reactions are usually exergonic reactions capable of releasing large amounts of energy

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3. Anabolic reactions (synthesis reactions) occur when a. The general chemical notation for this reaction is A + B  AB. b. These reactions are endergonic, fueled by chemical energy. D. Reaction Rates and Enzymes (Figure 2.8, 2.9, 2.10): For a reaction to occur, atoms must collide with enough energy overcome the repulsion of their electrons. This energy required for all chemical reactions is called the (Ea). The following factors increase the reaction rate by either reducing the activation energy or increasing the likelihood of strong collisions between reactants: concentration, temperature, reactant properties, and the presence or absence of a catalyst. Describe each of these factors that influence reaction rate. 1. Concentration:

2. Temperature:

3. Particle size and phase:

a. b.

4. Catalysts:

5. Most enzymes are macromolecule proteins. Summarize the properties of enzymes. a. b.

c.

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d. 6. Induced-fit mechanism:

(Figure 2.10) Module 2.4 Inorganic Compounds: Water, Acids, Bases, and Salts Bonds (Figures 2.11–2.13) A. Biochemistry, grouped into inorganic and organic compounds, is the chemistry of life. 1. Inorganic compounds generally do not contain

bonded to

hydrogen, and include water, acids, bases, and salts. 2. Organic compounds are defined as those that do contain bonded to hydrogen. B. Water (Figure 2.11): Water (H2O) makes up 60-80% of the mass of the human body and has the several key properties vital to our existence. 1. Summarize the properties of water. a.

b.

c.

d. 2. Water serves as the body’s primary solvent and is often called the universal solvent because so many solutes will dissolve in it entirely, or to some degree (Figure 2.11). a. Water is a polar covalent molecule where the oxygen pole is partially negative (d-) and the hydrogen pole is partially positive (d+), which allows the molecule to interact with certain solutes, surround them, and keep them apart.

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b. Water is only able to dissolve solutes that are or those with fully or partially charged ends. “Like dissolves like”: water generally dissolves ionic and polar covalent solutes. c. Solutes that do not have full or partially charged ends are known as and do not dissolve in water. This group includes uncharged nonpolar covalent molecules such as oils and fats. C. Acids and Bases (Figures 2.12, 2.13): The study of acids and bases is really the study of the hydrogen ion (H+) A water molecule in a solution may break apart or dissociate into a positively charged hydrogen ion and a negatively charged hydroxide ion (OH-). Acids and bases are defined in the following way according to their behavior with respect to hydrogen ions: 1. Define acid. (Figure 2.12b) 2. Define base. (Figure 2.12c) 3. The pH scale, ranging from

-

, is a simple way of representing the

hydrogen ion concentration of a solution and literally stands for negative logarithm of the hydrogen ion concentration, or -log [H+] (Figure 2.13). a. When the pH = 7 the solution is

, where the number of

hydrogen ions and base ions are equal. b. A solution with a pH less than 7 is

, where hydrogen

ions outnumber base ions. c. A solution with a pH greater than 7 is

or

, where base ions outnumber hydrogen ions. d. Most body fluids are slightly basic pH: blood,

-

; inside cells 7.2. 4. What is a buffer?

D. Salts and Electrolytes: A

refers any metal cation and nonmetal anion

held together by ionic bonds. Salts can dissolve in water to form cations and

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anions called

, which are capable of conducting an electrical

current. Module 2.5 Organic Compounds: Carbohydrates, Lipids, Proteins, and Nucleotides (Figures 2.14–2.26; Table 2.3) A. Monomers and Polymers: Each type of organic compound in the body, carbohydrate, lipid, protein, or nucleic acid, has its own monomer or single subunit and the corresponding polymer built from those subunits. 1. Describe dehydration synthesis.

2. Describe hydrolysis.

B. Carbohydrates (Figures 2.14, 2.15, 2.16; Table 2.3): Carbohydrates, composed of carbon, hydrogen, and oxygen, function primarily as fuel in the body with some limited structural roles. 1. Monosaccharides have from 3 to 7 carbons and are the monomers from which all carbohydrates are made. List examples of the most abundant monosaccharides in the body. ,

, , and

, (Figure 2.14).

2. Disaccharides are formed by union of two monosaccharides by dehydration synthesis (Figure 2.15). 3. Polysaccharides consist of many monosaccharides joined to one another by dehydration synthesis reactions (Figure 2.16). a. Glycogen is the storage polymer for

found mostly in

skeletal muscle and liver cells. b. Some polysaccharides are found covalently bound to either proteins or lipids resulting in the following compounds glycoproteins and glycolipids, which have various functions in the body. C. Lipids (Figures 2.17, 2.18, 2.19, 2.20; Table 2.3): The lipids, a group of nonpolar hydrophobic molecules composed primarily of carbon and hydrogen, include fats and oils.

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1. Fatty acids are the basic lipid monomers consisting of 4 to 20 carbon atoms which may have none, one, or more double bonds between the carbons in the hydrocarbon chain (Figure 2.17). a. Saturated fatty acids,

at room temperature, have

double bonds between carbon atoms so the carbons are “saturated” with the maximum number of hydrogen atoms (Figure 2.17a). b. Monounsaturated fatty acids, generally temperature, have

at room

double bond between two carbons in the

hydrocarbon chain (Figure 2.17b). c. Polyunsaturated fatty acids, liquid at room temperature, have separate double bonds between carbons in the hydrocarbon chain (Figure 2.17c). 2. Three fatty acids linked by dehydration synthesis to a modified 3-carbon carbohydrate, glycerol, form a triglyceride, the storage polymer for fatty acids also called a neutral fat (Figure 2.18). 3. Phospholipids are composed of a glycerol backbone, two fatty acid “tails”, and one phosphate group “head” in place of the third fatty acid (Figure 2.19). a. A molecule with a polar group, the phosphate head, and a nonpolar group, the fatty acid tail, is called amphiphilic. b. This amphiphilic nature makes phospholipids vital to the structure of cell membranes. 4. Steroids are nonpolar and share a four-ring hydrocarbon structure called the steroid nucleus.

is the steroid that forms the basis for all

the other steroids in the body (Figure 2.20). D. Proteins (Figures 2.21, 2.22, 2.23; Table 2.3): Proteins are macromolecules that are involved in movement, function as enzymes, play structural roles, function in the body’s defenses, and can be used as fuel. 1. Twenty different amino acids, the monomers of all proteins, can be linked by bonds into polypeptides (Figure 2.21).

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2. Peptides are formed when two or more amino acids are linked together by peptide bonds as a result of dehydration synthesis (Figure 2.22). a.

consist of two amino acids, tripeptides have three amino acids, andcontain 10 or more amino acids.

b.

consist of one or more polypeptide chains folded into distinct structures that must be maintained to be functional.

3. There are two basic types of proteins classified according to their structure: fibrous and globular (Figure 2.23). a. Describe fibrous proteins.

b. Describe globular proteins.

4. The complex structure of a complete protein is divided into four levels (Figure 2.23). Describe each level of protein structure. a. (Figure 2.23a) b.

(Figure 2.23b) c.

(Figure 2.23c) d.

(Figure 2.23d) 5. Explain protein denaturation.

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E. Nucleotides and Nucleic Acids (Figures 2.24, 2.25, 2.26; Table 2.3): Nucleotides are the monomers of nucleic acids, so named because of their abundance in the nucleus of cells, and make up our genetic material. 1. The nucleotide structure is composed of 3 parts. List the three parts (Figure 2.24a). a. b.

c. 2. There are two types of nitrogenous bases: purines and pyrimidines (Figure 2.24b). a. Purines, a double-ringed molecule, include

(A) and

(G). b. Pyrimidines, a single-ringed molecule, include (U), and

(C),

(T).

3. Adenosine triphosphate (ATP), adenine attached to ribose and three phosphate groups, is the main source of chemical energy in the body (Figure 2.25a). a. ATP is synthesized from

and a using energy from the oxidation of fuels

such as glucose. It is the potential energy in this “high-energy” bond that can be released to as kinetic energy to do work (Figure 2.25b). b. The production of large quantities of ATP requires oxygen, which is the main reason why we breathe air. 4. Deoxyribonucleic acid (DNA) and ribonucleic acid (RNA) are the two main nucleic acids that together are responsible for the storage and execution of the genetic code (Figure 2.26).

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5. DNA, an extremely large molecule found in the nucleus of the cell, is composed of two long chains that twist around each other to form a double helix (Figure 2.26a). Other structural features of DNA include: a. DNA contains the pentose sugar

, so called

because it is missing an oxygen-containing group found in ribose. b. DNA contains the following bases:

,

, and

,

.

c. The two stands of the double helix are held together by hydrogen bonding between the bases of each strand. d. DNA exhibits complementary base pairing where the purine always pairs with the pyrimidine T and the purine G always pairs with the pyrimidine

.

e. A = T (where = denotes 2 hydrogen bonds) and C ≡ G (where ≡ denotes 3 hydrogen bonds). This arrangement is allowed because each base faces the inside of the double helix as they run in opposite directions. f. DNA contains the genes that provide the recipe or code for protein synthesis, the process of making every protein in the body. 6. RNA, a single strand of nucleotides, can move between the nucleus of a cell and its cytosol and is critical to the making of proteins (Figure 2.26b). a. RNA contains the pentose sugar

.

b. RNA contains the pyrimidine uracil instead of thymine, which still pairs with adenine (A = U). c. RNA copies the recipe for a specific protein found in a gene on DNA, a process called

.

d. RNA is free to exit the nucleus to a location where protein synthesis occurs then proceeds to direct the making of the protein from the recipe, a process called

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.

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