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PowerPoint® Lecture Slides prepared by Janice Meeking, Mount Royal College

C H A P T E R

Copyright © 2010 Pearson Education, Inc.

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Chemistry Comes Alive: Part A


Objectives: Chapter Two

• Differentiate between matter and energy and between potential energy and kinetic energy.

• Define chemical element and list the four elements that form the bulk of body matter.

• Define molecule, and distinguish between a compound and a molecule.

• Differentiate among ionic, covalent, and hydrogen bonds.

• Compare and contrast polar and nonpolar bonds

• Define the three major types of chemical reactions: synthesis, decomposition, and exchange. Comment on the nature of oxidation-reduction reactions and their importance.

• Explain the importance of water and salts to body homeostasis.


Objectives Continued

• Define acid and base, and explain the concept of pH.

• Describe and compare the building blocks, general structures, and biological functions of carbohydrates and lipids.

• Explain the role of dehydration synthesis and hydrolysis in the formation and breakdown of organic molecules.

• Describe how enzymes function.

• Explain the role of ATP in cell metabolism.


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


Energy

• Capacity to do work or put matter into motion

• Types of energy:

• Kinetic—energy in action

• Potential—stored (inactive) energy


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)


Energy Form Conversions

• Energy may be converted from one form to another, but is considered “inefficient”

• Conversion is inefficient because some energy is given off as heat and may not be used. Heat is the lowest and most disorganized form of energy.


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


Major Elements of the Human Body

• Oxygen (O)

• Carbon (C)

• Hydrogen (H)

• Nitrogen (N)

About 96% of body mass


Lesser Elements of the Human Body

• About 3.9% of body mass:

• Calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe)


Trace Elements of the Human Body

• < 0.01% of body mass:

• Part of enzymes, e.g., chromium (Cr), manganese (Mn), and zinc (Zn)


Atomic Structure

• Determined by numbers of subatomic particles

• Nucleus consists of neutrons and protons


Atomic Structure

• Neutrons

• No charge

• Mass = 1 atomic mass unit (amu)

• Protons

• Positive charge

• Mass = 1 amu


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



• 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



• Atomic weight = average of mass numbers of all isotopes

Copyright © 2010 Pearson Education, Inc. Figure 2.3

Proton

Neutron

Electron

Deuterium (2H)(1p+; 1n0; 1e–)

Tritium (3H)(1p+; 2n0; 1e–)

Hydrogen (1H)(1p+; 0n0; 1e–)


Radioisotopes

• Spontaneous decay (radioactivity)

• Similar chemistry to stable isotopes

• Can be detected with scanners


Radioisotopes

• Valuable tools for biological research and medicine

• Cause damage to living tissue:

• Useful against localized cancers

• Radon from uranium decay causes lung cancer


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 atoms bonded together (e.g., C6H12O6)


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


Concentration of Solutions

• Expressed as

• Percent, or parts per 100 parts

• Milligrams per deciliter (mg/dl)

• Molarity, or moles per liter (M)

• 1 mole = the atomic weight of an element or molecular weight (sum of atomic weights) of a compound in grams

• 1 mole of any substance contains 6.02 1023 molecules (Avogadro’s number)


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


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)


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


Types of Chemical Bonds

• Ionic

• Covalent

• Hydrogen


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

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.


Formation of an Ionic Bond

• Ionic compounds form crystals instead of individual molecules

• NaCl (sodium chloride)

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


or

Oxygenatom

Oxygenatom

Molecule ofoxygen gas (O2)

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

oxygen atoms share two electron pairs.


+


+ or

Nitrogenatom

Nitrogenatom

Molecule ofnitrogen gas (N2)

Structuralformulashowstriple bond.(c) Formation of a triple covalent bond: Two

nitrogen atoms share three electron pairs.



Covalent Bonds

• Sharing of electrons may be equal or unequal

• Equal sharing produces electrically balanced nonpolar molecules

• CO2


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

• Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule

• Common between dipoles such as water

• Also act as intramolecular bonds, holding a large molecule in a three-dimensional shape


(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


(b) A water strider can walk on a pond because of the highsurface tension of water, a result of the combinedstrength of its hydrogen bonds.


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


Examples of Chemical Equations

H + H H2 (hydrogen gas)

4H + C CH4 (methane)

(reactants) (product)


Patterns of Chemical Reactions

• Synthesis (combination) reactions

• Decomposition reactions

• Exchange reactions


Synthesis Reactions

• A + B AB

• Always involve bond formation

• Anabolic


ExampleAmino acids are joined together toform a protein molecule.

(a) Synthesis reactions

Smaller particles are bondedtogether to form larger,

more complex molecules.

Amino acidmolecules

Proteinmolecule


Decomposition Reactions

• AB A + B

• Reverse synthesis reactions

• Involve breaking of bonds

• Catabolic


ExampleGlycogen is broken down to releaseglucose units.

Bonds are broken in largermolecules, resulting in smaller,

less complex molecules.

(b) Decomposition reactions

Glucosemolecules

Glycogen


Exchange Reactions

• AB + C AC + B

• Also called displacement reactions

• Bonds are both made and broken


ExampleATP transfers its terminal phosphategroup to glucose to form glucose-phosphate.

Bonds are both made and broken(also called displacement reactions).

(c) Exchange reactions

Glucose Adenosine triphosphate (ATP)

Adenosine diphosphate (ADP)Glucosephosphate

+

+


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

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