Copyright © 2010 Pearson Education, Inc. TEST REVIEW – EXAM 2 CHAPTER 2.

Copyright © 2010 Pearson Education, Inc.

TEST REVIEW – EXAM 2

CHAPTER 2


Terms to Know

• Molecule

• Atom

• Element

• Neutron

• Cation

• Nonpolar covalent bond

• Hydrogen bond

• Ionic bond

• Polar covalent bond

• Compound

• Mixture

• Suspensions

• Solutions

• Colloids

• Chemical energy

• Radiant energy

• Electrical energy

• Mechanical energy




Matter

• Anything that has mass and occupies space

• States of matter:

1. Solid—definite shape and volume

2. Liquid—definite volume, changeable shape

3. Gas—changeable shape and volume


Mass and Weight

• the mass of an object is a fundamental property of the object

• a numerical measure of its inertia

• measure of the amount of matter in the object.

• definitions of mass often seem circular because it is such a fundamental quantity that it is hard to define in terms of something else

• the usual symbol for mass is m and its SI unit is the kilogram

• the weight of an object is the force of gravity on the object (w = mg)

6


A. Weight

B. Matter

C. Mass

D. Energy

1. Can be measured only by its effects on matter.

2. Anything that occupies space and has mass.

3. Although a man who weighs 175 pounds on Earth would be lighter on the moon and heavier on Jupiter, his ________ would not be different.

4. Is a function of and varies with gravity.


Energy Concepts• What is energy?

• The capacity to perform work

• What is the difference between potential and kinetic energy?

• Stored vs. motion

• Energy is neither created nor destroyed but…

• Converted from one form to another

• This property is called the conservation of energy

• What is the usual way in which energy is “lost?”

• Through heat

• What type of energy is heat?

• Kinetic due to random motion of atoms

• Heat is generated by friction (in this example between atoms and air)

• Heat is highly __________ energy and highest amount of _________.

• Disordered, entropy

• Chemical energy is a form of ____________ energy.

• Potential

• What is the primary form of chemical energy in living organisms?

• ATP

• What is cellular respiration? What are the byproducts?

• Conversion of glucose into ATP through reduction of oxygen forming water and carbon dioxide


Forms of Energy

• Chemical energy — stored in bonds of chemical substances

• Electrical energy — results from movement of charged particles

• Mechanical energy — directly involved in moving matter

• Radiant or electromagnetic energy — exhibits wavelike properties (i.e., visible light, ultraviolet light, and X-rays)


Composition of Matter

• Elements

• Cannot be broken down by ordinary chemical means

• Each has unique properties:

• Physical properties

• Are detectable with our senses or are measurable

• Chemical properties

• How atoms interact (bond) with one another


Composition of Matter

• Atoms

• Unique building blocks for each element

• Atomic symbol: one- or two-letter chemical shorthand for each element


Atomic Structure

• Neutrons

• No charge

• Mass = 1 atomic mass unit (amu)

• Protons

• Positive charge

• Mass = 1 amu


Atomic Structure

• Determined by numbers of subatomic particles

• Nucleus consists of neutrons and protons


Atomic Structure

• Electrons

• Orbit nucleus

• Equal in number to protons in atom

• Negative charge

• 1/2000 the mass of a proton (0 amu)


Identifying Elements

• Atoms of different elements contain different numbers of subatomic particles

• Compare hydrogen, helium and lithium (next slide)

Copyright © 2010 Pearson Education, Inc. Figure 2.2

Proton

Neutron

Electron

Helium (He)(2p+; 2n0; 2e–)

Lithium (Li)(3p+; 4n0; 3e–)

Hydrogen (H)(1p+; 0n0; 1e–)



• Atomic number = number of protons in nucleus



• Mass number = mass of the protons and neutrons

• Mass numbers of atoms of an element are not all identical

• Isotopes are structural variations of elements that differ in the number of neutrons they contain


Radioisotopes

• Valuable tools for biological research and medicine

• Cause damage to living tissue:

• Useful against localized cancers

• Radon from uranium decay causes lung cancer


Proton

Neutron

Electron

Deuterium (2H)(1p+; 1n0; 1e–)

Tritium (3H)(1p+; 2n0; 1e–)

Hydrogen (1H)(1p+; 0n0; 1e–)


Chemically Inert Elements

• Stable and unreactive

• Outermost energy level fully occupied or contains eight electrons

Copyright © 2010 Pearson Education, Inc. Figure 2.5a

Helium (He)(2p+; 2n0; 2e–)

Neon (Ne)(10p+; 10n0; 10e–)

2e 2e8e

(a) Chemically inert elements

Outermost energy level (valence shell) complete


Chemically Reactive Elements

• Outermost energy level not fully occupied by electrons

• Tend to gain, lose, or share electrons (form bonds) with other atoms to achieve stability

Copyright © 2010 Pearson Education, Inc. Figure 2.5b

2e4e

2e8e

1e

(b) Chemically reactive elementsOutermost energy level (valence shell) incomplete

Hydrogen (H)(1p+; 0n0; 1e–)

Carbon (C)(6p+; 6n0; 6e–)

1e

Oxygen (O)(8p+; 8n0; 8e–) Sodium (Na)

(11p+; 12n0; 11e–)

2e6e


Molecules and Compounds

• Most atoms combine chemically with other atoms to form molecules and compounds

• Molecule — two or more atoms bonded together (e.g., H2 or C6H12O6)

• Compound — two or more different kinds of elements bonded together (e.g., C6H12O6)


Mixtures vs. Compounds

• Mixtures

• No chemical bonding between components

• Can be separated physically, such as by straining or filtering

• Heterogeneous or homogeneous

• Compounds

• Can be separated only by breaking bonds

• All are homogeneous


Mixtures

• Most matter exists as mixtures

• Two or more components physically intermixed

• Three types of mixtures

• Solutions

• Colloids

• Suspensions


Solutions

• Homogeneous mixtures

• Usually transparent, e.g., atmospheric air or seawater

• Solvent

• Present in greatest amount, usually a liquid

• Solute(s)

• Present in smaller amounts


Colloids and Suspensions

• Colloids (emulsions)

• Heterogeneous translucent mixtures, e.g., cytosol

• Large solute particles that do not settle out

• Undergo sol-gel transformations

• Suspensions:

• Heterogeneous mixtures (blood)

• Large visible solutes tend to settle out


Solution

Soluteparticles

Soluteparticles

Soluteparticles

Solute particles are verytiny, do not settle out or

scatter light.

ColloidSolute particles are larger

than in a solution and scatterlight; do not settle out.

SuspensionSolute particles are very

large, settle out, and mayscatter light.

ExampleMineral water

ExampleGelatin

ExampleBlood


Mixtures vs. Compounds

• Mixtures

• No chemical bonding between components

• Can be separated physically, such as by straining or filtering

• Heterogeneous or homogeneous

• Compounds

• Can be separated only by breaking bonds

• All are homogeneous


• Heterogeneous, will not settle.

• Heterogeneous, will settle.

• Homogeneous, will not settle.

•Will not scatter light.

A)Suspensions B) Solutions C) Colloids


• Nonpolar covalent bond

• Hydrogen bond

• Ionic bond

• Polar covalent bond


Chemical Bonds

• Electrons occupy up to seven electron shells (energy levels) around nucleus

• Octet rule: Except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their outermost energy level (valence shell)



Ionic Bonds

• Ions are formed by transfer of valence shell electrons between atoms

• Anions (– charge) have gained one or more electrons

• Cations (+ charge) have lost one or more electrons

• Attraction of opposite charges results in an ionic bond


Formation of an Ionic Bond

• Ionic compounds form crystals instead of individual molecules

• NaCl (sodium chloride)

Copyright © 2010 Pearson Education, Inc. Figure 2.6a-b

Sodium atom (Na)(11p+; 12n0; 11e–)

Chlorine atom (Cl)(17p+; 18n0; 17e–)

Sodium ion (Na+) Chloride ion (Cl–)

Sodium chloride (NaCl)

+ –

(a) Sodium gains stability by losing one electron, and chlorine becomes stable by gaining one electron.

(b) After electron transfer, the oppositely charged ions formed attract each other.

Copyright © 2010 Pearson Education, Inc. Figure 2.6c

CI–

Na+

(c) Large numbers of Na+ and Cl– ions associate to form salt (NaCl) crystals.


Covalent Bonds

• Formed by sharing of two or more valence shell electrons

• Allows each atom to fill its valence shell at least part of the time


+

Hydrogenatoms

Carbonatom

Molecule ofmethane gas (CH4)

Structuralformulashows singlebonds.

(a) Formation of four single covalent bonds: carbon shares four electron pairs with four hydrogen atoms.

or

Resulting moleculesReacting atoms


Covalent Bonds

• Sharing of electrons may be equal or unequal

• Equal sharing produces electrically balanced nonpolar molecules

• CO2


or

Oxygenatom

Oxygenatom

Molecule ofoxygen gas (O2)

Structuralformulashowsdouble bond.(b) Formation of a double covalent bond: Two

oxygen atoms share two electron pairs.

Resulting moleculesReacting atoms

+



Covalent Bonds

• Unequal sharing by atoms with different electron-attracting abilities produces polar molecules

• H2O

• Atoms with six or seven valence shell electrons are electronegative, e.g., oxygen

• Atoms with one or two valence shell electrons are electropositive, e.g., sodium



Hydrogen bonds• The bonds of a water molecule represent ________ _______ type of bond. Also known as a ________.

• Polar covalent, dipole

• Oxygen has a greater affinity for the electrons and is therefore more _____________. Whereas, hydrogen has a lesser attraction for electrons is more _____________.

• Electronegative, electropositive

• The oxygen end of the molecule is therefore slightly more _________ and the hydrogen ends are slightly more _________.

• Negative, positive

• The attraction between the negative oxygen end of one water compound to the positive hydrogen end of another water represents a ___________ bond.

• Hydrogen

• Hydrogen bonds are strong bonds. (T/F)

• False

• They are easily broken

• Hydrogen bonds may inter- or intramolecular. (T/F)

• True

• The unique properties of water are attributable to hydrogen bonds. Some of the properties include….

• Cohesion, high boiling point, why ice floats, high heat of vaporization, high heat capacity


(a) The slightly positive ends (+) of the watermolecules become aligned with the slightlynegative ends (–) of other water molecules.

+

–

–

–– –

+

+

+

+

+

Hydrogen bond(indicated bydotted line)

Figure 2.10a


Chemical Reactions

• Occur when chemical bonds are formed, rearranged, or broken

• Represented as chemical equations

• Chemical equations contain:

• Molecular formula for each reactant and product

• Relative amounts of reactants and products, which should balance


Patterns of Chemical Reactions

• Synthesis (combination) reactions

• Decomposition reactions

• Exchange reactions


Synthesis Reactions

• A + B AB

• Always involve bond formation

• Anabolic


• What is dehydration synthesis?

• Removal of a water molecule to form a new covalent bond

• What is hydrolysis?

• The addition of a water molecule to break a covalent bond

• What is anabolism?

• Forming new bonds to build something bigger. Requires energy (endergonic)

• What is catabolism?

• Breaking bonds to make something smaller. Large molecules down to subunits.

• Releases energy (exergonic).

Dehydration Synthesis and Hydrolysis


Decomposition Reactions

• AB A + B

• Reverse synthesis reactions

• Involve breaking of bonds

• Catabolic


Oxidation-Reduction (Redox) Reactions

• Decomposition reactions: Reactions in which fuel is broken down for energy

• Also called exchange reactions because electrons are exchanged or shared differently

• Electron donors lose electrons and are oxidized

• Electron acceptors receive electrons and become reduced


Chemical Reactions

• All chemical reactions are either exergonic or endergonic

• Exergonic reactions — release energy

• Catabolic reactions

• Endergonic reactions — products contain more potential energy than did reactants

• Anabolic reactions


Chemical Reactions

• All chemical reactions are theoretically reversible

• A + B AB

• AB A + B

• Chemical equilibrium occurs if neither a forward nor reverse reaction is dominant

• Many biological reactions are essentially irreversible due to

• Energy requirements

• Removal of products


Rate of Chemical Reactions

• Rate of reaction is influenced by:

• temperature rate

• particle size rate

• concentration of reactant rate

• Catalysts: rate without being chemically changed

• Enzymes are biological catalysts


Classes of Compounds

• Inorganic compounds

• Water, salts, and many acids and bases

• Do not contain carbon

• Organic compounds

• Carbohydrates, fats, proteins, and nucleic acids

• Contain carbon, usually large, and are covalently bonded


Water

• 60%–80% of the volume of living cells

• Most important inorganic compound in living organisms because of its properties


Properties of Water

• High heat capacity

• Absorbs and releases heat with little temperature change

• Prevents sudden changes in temperature

• High heat of vaporization

• Evaporation requires large amounts of heat

• Useful cooling mechanism


Properties of Water

• Polar solvent properties

• Dissolves and dissociates ionic substances

• Forms hydration layers around large charged molecules, e.g., proteins (colloid formation)

• Body’s major transport medium


Water molecule

Ions in solutionSalt crystal

–

+

+


Properties of Water

• Reactivity

• A necessary part of hydrolysis and dehydration synthesis reactions

• Cushioning

• Protects certain organs from physical trauma, e.g., cerebrospinal fluid


Salts

• Ionic compounds that dissociate in water

• Contain cations other than H+ and anions other than OH–

• Ions (electrolytes) conduct electrical currents in solution

• Ions play specialized roles in body functions (e.g., sodium, potassium, calcium, and iron)


Acids and Bases

• Both are electrolytes

• Acids are proton (hydrogen ion) donors (release H+ in solution)

• HCl H+ + Cl–


Acids and Bases

• Bases are proton acceptors (take up H+ from solution)

• NaOH Na+ + OH–

• OH– accepts an available proton (H+)

• OH– + H+ H2O

• Bicarbonate ion (HCO3–) and ammonia

(NH3) are important bases in the body because of buffering properties


Acid-Base Concentration

• Acid solutions contain [H+]

• As [H+] increases, acidity increases, pH decreases

• Alkaline solutions contain bases (e.g., OH–)

• As [H+] decreases (or as [OH–] increases), alkalinity increases, pH increases


pH: Acid-Base Concentration

• pH = the negative logarithm of [H+] in moles per liter

• Neutral solutions:

• Pure water is pH neutral (contains equal numbers of H+ and OH–)

• pH of pure water = pH 7: [H+] = 10 –7 M

• All neutral solutions are pH 7


pH: Acid-Base Concentration

• Acidic solutions

• [H+], pH

• Acidic pH: 0–6.99

• pH scale is logarithmic: a pH 5 solution has 10 times more H+ than a pH 6 solution

• Alkaline solutions

• [H+], pH

• Alkaline (basic) pH: 7.01–14


Concentration(moles/liter)

[OH–]

100 10–14

10–1 10–13

10–2 10–12

10–3 10–11

10–4 10–10

10–5 10–9

10–6 10–8

10–7 10–7

10–8 10–6

10–9 10–5

10–10 10–4

10–11 10–3

10–12 10–2

10–13 10–1

[H+] pHExamples

1M Sodiumhydroxide (pH=14)

Oven cleaner, lye(pH=13.5)

Household ammonia(pH=10.5–11.5)

Neutral

Household bleach(pH=9.5)

Egg white (pH=8)

Blood (pH=7.4)

Milk (pH=6.3–6.6)

Black coffee (pH=5)

Wine (pH=2.5–3.5)

Lemon juice; gastricjuice (pH=2)

1M Hydrochloricacid (pH=0)10–14 100

14

13

12

11

10

9

8

7

6

5

4

3

2

1

0


Acid-Base Homeostasis

• pH change interferes with cell function and may damage living tissue

• Slight change in pH can be fatal

• pH is regulated by kidneys, lungs, and buffers


Buffers

• Mixture of compounds that resist pH changes

• Convert strong (completely dissociated) acids or bases into weak (slightly dissociated) ones

• Carbonic acid-bicarbonate system


Organic Compounds

• Contain carbon (except CO2 and CO, which are inorganic)

• Unique to living systems

• Include carbohydrates, lipids, proteins, and nucleic acids


Organic Compounds

• Many are polymers — chains of similar units (monomers or building blocks)

• Synthesized by dehydration synthesis

• Broken down by hydrolysis reactions

• How are polymers formed?

• By dehydration synthesis

• What reactions break down polymers into monomers?

• By hydrolysis

• What molecule is essential to this process?

• H2O


+

Glucose Fructose

Water isreleased

Monomers linked by covalent bond

Monomers linked by covalent bond

Water isconsumed

Sucrose

(a) Dehydration synthesis

Monomers are joined by removal of OH from one monomerand removal of H from the other at the site of bond formation.

+

(b) Hydrolysis

Monomers are released by the addition of a water molecule, adding OH to one monomer and H to the other.

(c) Example reactions

Dehydration synthesis of sucrose and its breakdown by hydrolysis

Monomer 1 Monomer 2

Monomer 1 Monomer 2

+


Carbohydrates

• Sugars and starches

• Contain C, H, and O [(CH20)n]

• Three classes

• Monosaccharides

• Disaccharides

• Polysaccharides


Carbohydrates

• Functions

• Major source of cellular fuel (e.g., glucose)

• Structural molecules (e.g., ribose sugar in RNA)


Monosaccharides

• Simple sugars containing three to seven C atoms

• (CH20)n n = 3 – 7

• C3H6O3

• C6H12O6


ExampleHexose sugars (the hexoses shown

here are isomers)

ExamplePentose sugars

Glucose Fructose Galactose Deoxyribose Ribose

(a) MonosaccharidesMonomers of carbohydrates


ExampleThis polysaccharide is a simplified representation of

glycogen, a polysaccharide formed from glucose units.

(c) PolysaccharidesLong branching chains (polymers) of linked monosaccharides

Glycogen

• Glycogen is animals main storage form of glucose. Found in high concentrations in the liver and muscles.

• Starch is plants main storage form of glucose.

• Cellulose is a key structural molecule in plants. Not digestible by humans.


Lipids

• Contain C, H, O (less than in carbohydrates), and sometimes P

• Insoluble in water

• Main types:

• Neutral fats or triglycerides

• Phospholipids

• Steroids

• Eicosanoids


Triglycerides

• Neutral fats — solid fats and liquid oils

• Composed of three fatty acids bonded to a glycerol molecule

• Main functions

• Energy storage

• Insulation

• Protection


Glycerol

+

3 fatty acid chains Triglyceride,or neutral fat

3 watermolecules

(a) Triglyceride formation

Three fatty acid chains are bound to glycerol bydehydration synthesis


Saturation of Fatty Acids

• Saturated fatty acids

• Single bonds between C atoms; maximum number of H

• Solid animal fats, e.g., butter

• Unsaturated fatty acids

• One or more double bonds between C atoms

• Reduced number of H atoms

• Plant oils, e.g., olive oil


Phospholipids

• Modified triglycerides:

• Glycerol + two fatty acids and a phosphorus (P)-containing group

• “Head” and “tail” regions have different properties (amphipathic)

• Hydrophilic head

• Hydrophobic tail

• Important in cell membrane structure


Phosphorus-containing

group (polar“head”)

ExamplePhosphatidylcholine

Glycerolbackbone

2 fatty acid chains(nonpolar “tail”)

Polar“head”

Nonpolar“tail”

(schematicphospholipid)

(b) “Typical” structure of a phospholipid molecule

Two fatty acid chains and a phosphorus-containing group areattached to the glycerol backbone.


Steroids

• Steroids — interlocking four-ring structure

• Cholesterol, vitamin D, steroid hormones, and bile salts


ExampleCholesterol (cholesterol is the

basis for all steroids formed in the body)

(c) Simplified structure of a steroid

Four interlocking hydrocarbon rings form a steroid.


Other Lipids in the Body

• Other fat-soluble vitamins

• Vitamins A, D, E, and K

• Lipoproteins

• Transport fats in the blood


Proteins

• Polymers of amino acids (20 types)

• Joined by peptide bonds

• Contain C, H, O, N, and sometimes S and P


(a) Generalized structure of all amino acids.

(b) Glycine is the simplest

amino acid.

(c) Aspartic acid (an acidic amino acid)

has an acid group (—COOH) in the

R group.

(d) Lysine (a basic amino acid) has an amine group

(–NH2) in the R group.

(e) Cysteine (a basic amino acid)

has a sulfhydryl (–SH) group in the R group, which suggests that

this amino acid is likely to participate in

intramolecular bonding.

Aminegroup

Acidgroup


Amino acid Amino acid Dipeptide

Dehydration synthesis:The acid group of one

amino acid is bonded to the amine group of the

next, with loss of a water molecule.

Hydrolysis: Peptide bonds linking amino

acids together are broken when water is

added to the bond.

+

Peptidebond


(a) Primary structure: The sequence of amino acids forms the polypeptide chain.

Amino acid Amino acid Amino acid Amino acid Amino acid


-Helix: The primary chain is coiledto form a spiral structure, which is

stabilized by hydrogen bonds.

-Sheet: The primary chain “zig-zags” backand forth forming a “pleated” sheet. Adjacentstrands are held together by hydrogen bonds.

(b) Secondary structure:The primary chain forms spirals (-helices) and sheets (-sheets).


Tertiary structure of prealbumin(transthyretin), a protein that

transports the thyroid hormonethyroxine in serum and cerebro-

spinal fluid.

(c) Tertiary structure: Superimposed on secondary structure. -Helices and/or -sheets are

folded up to form a compact globular molecule held together by intramolecular bonds.

Copyright © 2010 Pearson Education, Inc. Figure 2.19d

Quaternary structure ofa functional prealbuminmolecule. Two identical

prealbumin subunitsjoin head to tail to form

the dimer.

(d) Quaternary structure: Two or more polypeptide chains, each with its own tertiary structure,

combine to form a functional protein.


Fibrous and Globular Proteins

• Fibrous (structural) proteins

• Strandlike, water insoluble, and stable

• Examples: keratin, elastin, collagen, and certain contractile fibers


Fibrous and Globular Proteins

• Globular (functional) proteins

• Compact, spherical, water-soluble and sensitive to environmental changes

• Specific functional regions (active sites)

• Examples: antibodies, hormones, molecular chaperones, and enzymes


Protein Denaturation

• Shape change and disruption of active sites due to environmental changes (e.g., decreased pH or increased temperature)

• Reversible in most cases, if normal conditions are restored

• Irreversible if extreme changes damage the structure beyond repair (e.g., cooking an egg)


Activationenergy required

Less activationenergy required

WITHOUT ENZYME WITH ENZYME

Reactants

Product Product

Reactants

Enzymes

Copyright © 2010 Pearson Education, Inc. Figure 2.21, step 3

Substrates (S)e.g., amino acids

Enzyme (E)

Enzyme-substratecomplex (E-S)

Enzyme (E)

Product (P)e.g., dipeptide

Energy isabsorbed;

bond isformed.

Water isreleased.

Peptidebond

Substrates bindat active site.

Enzyme changesshape to holdsubstrates in

proper position.

Internalrearrangements

leading tocatalysis occur.

Product isreleased. Enzymereturns to original

shape and isavailable to catalyzeanother reaction.

Active site

+ H2O

1 23


Enzymes• What is an enzyme?

• Protein

• Biologic catalyst

• What is a catalyst

• Substance that speeds up a reaction

• What is Ea?

• Energy of activation

• Enzymes do what to a reaction?

• Lower energy of activation (heat, mechanical, chemical, etc)

• Speeds up rxn

• On what does an enzyme act?

• Its substrate

• Enzymes are __________ for their substrates?

• Specific


Nucleic Acids

• DNA and RNA

• Largest molecules in the body

• Contain C, O, H, N, and P

• Building block = nucleotide, composed of N-containing base, a pentose sugar, and a phosphate group


Deoxyribonucleic Acid (DNA)

• Four bases:

• adenine (A), guanine (G), cytosine (C), and thymine (T)

• Double-stranded helical molecule in the cell nucleus

• Provides instructions for protein synthesis

• Replicates before cell division, ensuring genetic continuity


Deoxyribosesugar

Phosphate

Sugar-phosphatebackbone

Adenine nucleotideHydrogen

bond

Thymine nucleotide

PhosphateSugar:

Deoxyribose PhosphateSugarThymine (T)Base:

Adenine (A)

Adenine (A)

Thymine (T)

Cytosine (C)

Guanine (G)

(b)

(a)

(c) Computer-generated image of a DNA molecule


Ribonucleic Acid (RNA)

• Four bases:

• adenine (A), guanine (G), cytosine (C), and uracil (U)

• Single-stranded molecule mostly active outside the nucleus

• Three varieties of RNA carry out the DNA orders for protein synthesis

• messenger RNA, transfer RNA, and ribosomal RNA


True / False

• Chemical properties are determined primarily by neutrons.

• No chemical bonding occurs between the components of a mixture.

• Buffers resist abrupt and large changes in the pH of the body by releasing or binding ions.

• All organic compounds contain carbon.

• A dipeptide can be broken into two amino acids by dehydration synthesis.

• The lower the pH, the higher the hydrogen ion concentration.

• Covalent bonds are generally less stable than ionic bonds.

• Hydrogen bonds are comparatively strong bonds.

• The fact that no chemical bonding occurs between the components of a mixture is the chief difference between mixtures and compounds.


True / False

• The pH of body fluids must remain fairly constant for the body to maintain homeostasis

• A charged particle is generally called an ion.

• Isotopes differ from each other only in the number of electrons contained.

• About 60% to 80% of the volume of most living cells consists of organic compounds.

• Lipids are a poor source of stored energy.

• Current information theorizes that omega-3 fatty acids decrease the risk of heart disease.

• Glucose is an example of a monosaccharide.

• A molecule consisting of one carbon atom and two oxygen atoms is correctly written as CO2.


Multiple Choice

• Choose the statement that is false or incorrect.

A) In chemical reactions, breaking old bonds requires energy and forming new bonds releases energy.

B) Exergonic reactions release more energy than they absorb.

C) A key feature of the body’s metabolism is the almost exclusive use of exergonic reactions by

the body.

D) Endergonic reactions absorb more energy than they release.

• A chemical reaction in which bonds are broken is usually associated with ________.

A) a synthesis B) the consumption of energy C) the release of energy D) forming a larger molecule

• What happens in redox reactions?

A) the reaction is always easily reversible B) the electron acceptor is oxidized

B) both decomposition and electron exchange occur D) the electron donor is reduced

• Choose the answer that best describes fibrous proteins

A) are usually called enzymes B) are very stable and insoluble in water

C) rarely exhibit secondary structure D) are cellular catalysts


Multiple Choice

• In liquid XYZ, you notice that light is scattered as it passes through. There is no precipitant in the

bottom of the beaker, though it has been sitting for several days. What type of liquid is this?

A) suspension B) solution C) mixture D) colloid

• Atom X has 17 protons. How many electrons are in its valence shell?

• A) 10 B) 5 C) 3 D) 7

• Which protein types are vitally important to cell function in all types of stressful circumstances?

• A) catalytic proteins B) molecular chaperones

• C) regulatory proteins D) structural proteins

• If atom X has an atomic number of 74 it would have which of the following?

A) 37 protons and 37 neutrons B) 37 protons and 37 electrons

C) 74 protons D) 37 electrons

• What does the formula C6H12O6 mean?

A) There are 6 calcium, 12 hydrogen, and 6 oxygen atoms.

• B) There are 12 hydrogen, 6 carbon, and 6 oxygen atoms.

• C) The substance is a colloid.

• D) The molecular weight is 24.

• Two good examples of a colloid would be Jell-O® and ________.

• A) cytosol B) blood C) urine D) toenails


Multiple Choice• An atom with a valence of 3 may have a total of ________ electrons.

A) 3 B) 8 C) 17 D) 13

• The chemical symbol O=O means ________.

• A) zero equals zero

• B) the atoms are double bonded

• C) both atoms are bonded and have zero electrons in the outer orbit

• D) this is an ionic bond with two shared electrons

• What is a dipole?

A) a type of reaction B) a type of bond

C) an organic molecule D) a polar molecule

• Amino acids joining together to make a peptide is a good example of a(n) ________ reaction.

A) decomposition B) reversible

C) exchange D) synthesis

• Which of the following is not considered a factor in influencing a reaction?

A) time B) concentration C) particle size D) temperature

• Which of the following is not an electrolyte?

A) NaOH B) HCl C) H2O D) Ca2CO3

Copyright © 2010 Pearson Education, Inc. TEST REVIEW – EXAM 2 CHAPTER 2.

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