Essentials of Biology Sylvia S. Mader Chapter 2 Lecture Outline Copyright © The McGraw-Hill...

Essentials of BiologySylvia S. Mader

Chapter 2Lecture Outline

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2.1 The Nature of MatterMatter

Anything that takes up space and has mass Can exist as a solid, liquid, or a gas Composed of elements

• Element cannot be broken down into another substance by

ordinary chemical means Pure substance consisting of one type of atom

• Only 92 naturally occurring elements• Four most common elements in living organisms

– CHON

Figure 2.1 Elements in

living organisms

calcium (Ca) 2%nitrogen (N) 5%

hydrogen (H) 10%oxygen

(O)65%

carbon (C)

18%

phosphorus (P) 1.1%

other elements,including sulfur 0.9%

© Tim Pannell/Corbis RF

Atomic structure Atomic theory

• Elements consist of atoms• Atom: smallest unit of an element that has the properties of the

element

Atomic symbols – one or two letters• H = ? C = ? S = ? Cl = ? Na = ?

Subatomic particles• Neutrons, n

no electrical charge, found in nucleus

• Protons, p+

positive charge, found in nucleus determines the ID of an atom

• Electrons , e- negative charge, outside of nucleus determines the chemical properties of an atom

Mass number • sum of protons and neutrons (electrons have nearly zero mass)

Figure 2.2 Two models of helium (He)

inside nucleus outside nucleus

= proton = neutron = electron

nucleus

a. b.

+

+

+

+

+ –

–

–

Atomic number All atoms of an element have this same number of

protons. Atoms are electrically neutral:

• How do the number of protons compare to the number of electrons?

• Periodic table Elements in order of their __________________ . Atoms arranged in periods (rows) and groups

(columns) Elements’ chemical and physical characteristics

recur in predictable manner• E.g. LiCl, NaCl, KCl; BeCl2, CaCl???

Periodic table of the elements

Location of....• Metals?• Nonmetals?

Figure 2.3 A portion of the periodic table

1

1

2

3

4

3

Li6.941

11

Na22.99

19

K39.10

4

Be9.012

12

Mg24.31

20

Ca40.08

5

B10.81

13

Al26.98

31

Ga69.72

6

C12.01

14

Si28.09

32

Ge72.59

7

N14.01

15

P30.97

33

As74.92

8

O16.00

16

S32.07

34

Se78.96

9

F19.00

17

Cl35.45

35

Br79.90

GroupsP

erio

ds

8

2

He4.003

10

Ne20.18

18

Ar39.95

36

Kr83.60

2 3 4 5 6 71.008

H1

Isotopes Atoms of the same element that differ in the

number of _________________ .

Isotopes have the same number of ________ but a different number of ___________ (different mass numbers)

Unstable isotopes may decay emitting radiation

Radioactive isotopes used…

• Can be used as tracer – PET scan

• Can cause damage to cells leading to cancer

• Can be used to sterilize medical equipment

Figure 2.4 PET Scans

thyroidgland

a.

b.

a: © Biomed Commun./Custom Medical Stock Photo; b(both):Courtesy National Institutes of Health

Figure 2.5 Chemotherapy using isotopes that emit high levels of radiation

a.b.

a: © Natasja Weitsz/Getty Images; b: © Geoff Tompkinson/SPL/Photo Researchers, Inc.

Arrangements of Electrons Located outside the nucleus of an atom in

specific electron shells (energy levels) Each shell contains a certain number of electrons For atoms up through number 20

• 2 electrons fill first shell.

• 8 electrons fill each additional shell.

Octet rule for valence shell• Valence shell: outermost shell

• Atoms are most stable with 8 valence electrons. Exceptions: atoms with the only one energy level

• Atoms can give up, accept, or share electrons to have 8.

The number of electrons in the valence shell determines the chemical properties of an atom

Atoms of the four elements most abundant in life

Electron

Firstelectron shell:can hold2 electrons

Outermostelectron shell:can hold8 electrons

Carbon (C)Atomic number = 6

Nitrogen (N)Atomic number = 7

Oxygen (O)Atomic number = 8

Hydrogen (H)Atomic number = 1

Orbital Diagrams of the First 18 Elements

2

8

8

1st Shell

2nd Shell

3rd Shell

Figure 2.6 Atoms of the six elements, CHNOPS


P

Phosphorus

O

Oxygen

N

Nitrogen

C

Carbon

H

Hydrogen

electron

1

H1

6

C12

7

N14

8

O16

15

P31

electronshell

nucleus

inner shell

outer(valence)shell

S

Sulfur16

S32

• Chemical reactions:

– Atoms give up or acquire electrons in order to complete their _____________ shells

– Result in atoms staying close together to form molecules

– Chemical bonds hold molecules together

• Ionic Bonds

• Covalent bonds

Chemical Bonding and Molecules

Types of chemical bonds Molecule – group of atoms chemically bonded

together• O2, H2O, C6H12O6, N2

Compound – substance containing atoms of more than one element

• H2O, C6H12O6

2 types of chemical bonds• Ionic bonds – giving up or accepting electrons

Form between metals and nonmetals

• Covalent bonds – sharing electrons Form between nonmetals

• When an atom loses or gains electrons, it becomes electrically charged. Why?

– Charged atoms are called ____________

– Ionic bonds are formed between oppositely charged ________________

Ionic Bonds: form between metals and nonmetals

Sodium atom (Na) Chlorine atom (Cl)

Completeouter shells

Sodium ion (Na) Chloride ion (Cl)

Sodium chloride (NaCl)Figure 2.7

(a) Hydrogen atom (H)

(c) Sodium atom (Na)

(b) Hydrogen ion (H+)

(d) Sodium ion (Na+)

1 electron

1 proton

No net electrical charge

11 electrons

11 protons

No net electrical charge

No electron

1 proton

10 electrons

11 protons

Fig. 2.03

Atoms: electrically neutral Ions: Electrically charged

Figure 2.7 Formation of sodium chloride

chlorine atom (Cl)sodium atom (Na)

Ionic bonding Forms when 2 atoms held together by the attraction between

opposite charges Sodium has _____electron in valence shell

• Usually _________________an electron to another atom

Chlorine has ____electrons in valence shell• Usually ________an electron from another atom

Na Cl

Figure 2.7 continued

Cl

chlorine atom (Cl)

Na

sodium atom (Na) +

Ions – charged atoms (No. P+ ≠ No. e-)

• Sodium ion, Na+

• 1 more _____________ than electrons

• Chloride ion, Cl-

• 1 more _____________ than protons

• Ionic compounds often called salts

• Covalent bonding 2 atoms ____________ electrons 2 hydrogen atoms can share electrons to fill

their first shell – orbitals overlap.

Structural formula – uses straight lines H-H• One line indicates 1 pair of shared electrons.

Molecular formula – shows number of atoms involved H2

Hydrogen gas (H2)

H H

Double covalent bond: sharing 2 pairs of electrons

Oxygen gas O2 or O=O

• Triple covalent bond – sharing __?__ pairs of electrons Nitrogen gas N2 or N≡N

• An atom may form bonds with more than one atom…

O O

Covalent bonding in water

Full shell with 8 electrons– Slightly negative

Hydrogen atoms with unfilled shells

Full shells with 2 electrons each

+Slightly positive

+

Covalent bond(shared pairof electron)

Oxygen atom with unfilled shell

Water molecule (H2O)

Covalent Bonds: form between nonmetallic atoms

• Cells constantly rearrange molecules by breaking existing chemical bonds and forming new ones

– Such changes in the chemical composition of matter are called chemical reactions

Chemical Reactions

Hydrogen gas Oxygen gas Water

Reactants Products

Reactants: on the left side of the equation – the starting

materials

Chemical Equations: symbolize chemical reactions

Products: on the right side of the equation – the ending materials (the stuff

produces)

Law of Conservation of Mass– Chemical reactions do not create or destroy matter—they only

rearrange it!

• Chemical reactions Reactants – molecules that participate in reaction

• Shown to left of arrow

Products – molecules formed by reactions• Shown to right of arrow

Equation is balanced if the same number of each type of atom occurs on both sides of arrow.

• An overall equation for photosynthesis

• Molecular formula for glucose

6 CO2

carbondioxide

6 H2Owater

C6H12O6

glucose+ 6 O2

oxygen+


C6H12O6

one molecule

indicates6 atomsof carbon

indicates6 atomsof oxygen

indicates12 atomsof hydrogen

• Life on Earth began in water and evolved there for 3 billion years

• The abundance of water is a major reason Earth is habitable

– Modern life still remains tied to water

– Your cells are composed of 70%–95% water

• Water has unique properties that make it a life-supporting substance.

• Properties stem from structure of molecule

2.2 Water’s Importance to Life

• The water molecule:

– two hydrogen atoms joined to one oxygen atom by single covalent bonds

The Structure of Water

H

O

H

• The electrons of the covalent bonds are shared unequally between oxygen and hydrogen

– unequal sharing of electrons makes water a polar molecule

– hydrogen atoms: partially positive ( ) Why?

– oxygen atom: partially negative ( -) Why?

()( )

( )

Water: a polar molecule

Figure 2.9 The structure of water

• The polarity of water results in weak electrical attractions between neighboring water molecules

– These interactions are called hydrogen bonds (b)

()

Hydrogen bond()

()()

()

()

()

()

The Structure of Water

Figure 2.9 The structure of water

• The polarity of water molecules and the hydrogen bonding that results explain most of water’s life-supporting properties

– Cohesion and adhesion

– High surface tension

– High heat capacity

– High heat of vaporization

– Varying density

Water’s Life-Supporting Properties

• A solution is a liquid consisting of two or more substances evenly mixed

Water as the Solvent of Life

– The dissolving agent is called the solvent

– The dissolved substance is called the solute

Ion in solutionSalt crystal

Dissolving of Sodium Chloride (NaCl) in Water

Salt

Water

Electricalattraction

Watermolecules(H2O)

Hydrogenbonds

Edge of onesalt crystal

Ionic bond

Water molecules dissolve NaCl,breaking ionic bond

• Water molecules stick together as a result of hydrogen bonding

– This is called cohesion

– Cohesion is vital for water transport in plants

The Cohesion of Water

Microscopic tubes

• Surface tension

– is the measure of how difficult it is to stretch or break the surface of a liquid

– Hydrogen bonds give water an unusually high surface tension

Figure 2.13

• Because of hydrogen bonding, water has a strong resistance to temperature change

• Water can absorb and store large amounts of heat while only changing a few degrees in temperature

– Earth’s Oceans cause temperatures to stay within limits that permit life

Water Moderates Temperature

• The density of ice is lower than liquid water

– This is why ice floats

Hydrogen bond

Liquid water

Hydrogen bondsconstantly break and re-form

Ice

Stable hydrogen bonds

• Water is a solvent. Due to polarity and H-bonding, water

dissolves many substances.

Hydrophilic – molecules attracted to water

Hydrophobic – molecules not attracted to water

Water causes NaCl to dissociate.

++ –

The salt NaCl dissociates in water.

Na+ Cl–

++–O

H HH H

O

• Cohesion Ability of water molecules to cling to each other due to

hydrogen bonding

• Adhesion Ability of water molecules to cling to other polar

surfaces

• Allows water to be excellent transport system in and outside of living organisms.

• Contributes to water transport in plants

Figure 2.10 Cohesion and adhesion of water

molecules


Wat

er in

wat

er c

olu

mn

Water evaporates, pullingthe water column fromthe roots to the leaves.

Water molecules clingtogether and adhere tosides of vessels in stemsand tree trunks.

H2O

Water enters a plantat root cells.

H2O

© Corbis RF

• Water has a high surface tension. Water molecules at the surface cling more

tightly to each other than to the air above. Mainly due to hydrogen bonding

• High heat capacity The many hydrogen bonds linking water

molecules allow water to absorb heat without greatly changing its temperature.

Temperature of water rises and falls slowly.

• High heat of vaporization Takes a great deal of energy to break H bonds for

evaporation Heat is dispelled as water evaporates.

Figure 2.12 Heat of vaporizationCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

b.

a.

a: © The McGraw-Hill Companies, Inc./Jill Braaten, photographer; b: © SuperStock RF

• Water is less dense than ice. Unlike other substances, water expands as it

freezes. Ice floats rather than sinks. It makes life possible in water. Ice acts as an insulator.

Figure 2.13 Properties of ice


a. Ice

b. Pond

ice layer

• When water molecules get cold, they move apart, forming ice

– A chunk of ice has fewer molecules than an equal volume of liquid water

• Since ice floats, ponds, lakes, and even the oceans do not freeze solid

– Marine life could not survive if bodies of water froze solid

The Biological Significance of Ice Floating

2.3 Acids and Bases• Water dissociates

into an equal number of hydrogen ions (H+) and hydroxide ions (OH-)

Figure 2.14 Dissociation of water molecules


OH H

OH H

OH H

OH–

OH–H+

H+


H – O – H

water

H+ + OH–

hydrogenion

hydroxideion

• Acids Common examples are

lemon juice, vinegar, tomatoes, and coffee.

Substances that dissociate in water, releasing H+ ions

Adding an acid to water increases the number of H+ ions.

HCl H+ + Cl-

Hydrochloric acid

Figure 2.15 Addition of hydrochloric acid (HCl)

OH H

OH H

OH H

OH–

Cl–

Cl–

OH–

HCl

H+

H+ H+

H+


• Bases Common bases are

ammonia and milk of magnesia.

Substances that either take up hydrogen ions or release hydroxide ions

Adding a base to water either increases the number of OH- ions or decreases the number of H+ ions.

NaOH Na+ + OH-

Sodium hydroxide

Figure 2.16 Addition of sodium hydroxide (NaOH), a base

O

H H

O

H H

O

H H

OH–

OH–

OH–

OH–

NaOH

Na+

Na+

H+

H+

• pH Mathematical way to indicate number of

hydrogen ions in solution pH scale ranges from 0 to 14

• pH below 7 acidic – more [H+] than [OH-]• pH above 7 basic – more [OH-] than [H+]• pH of 7 neutral – [H+] equal to [OH-]

• Acid A chemical compound that donates H+ ions to

solutions Taste sour

• Base A compound that ….

• accepts H+ ions and removes them from solution• Dissolves in water to produce hydroxide ions, OH• Taste bitter

Acids, Bases, and pH

Basicsolution

Neutralsolution

Acidicsolution

Oven cleaner

Household bleach

Household ammonia

Milk of magnesia

Seawater

Human bloodPure water

Urine

Tomato juice

Grapefruit juice

Lemon juice;gastric juice

• Acidic: pH _?_ 7

H+ _?_ OH-

• Basic: pH _?_ 7

H+ _?_ OH-

• Neutral: pH _?_ 7

H+ __?__ OH-

pH ScaleThe pH scale is used to describe the acidity of a

solution

Basicsolution

Neutralsolution

Acidicsolution

Oven cleaner

Household bleach

Household ammonia

Milk of magnesia

Seawater

Human bloodPure water

Urine

Tomato juice

Grapefruit juice

Lemon juice;gastric juice

• Acidic: pH < 7

H+ > OH-

• Basic: pH > 7

H+ < OH-

• Neutral: pH = 7

H+ = OH-

pH ScaleThe pH scale is used to describe the acidity of a

solution

• Buffer Chemical or combination of chemicals that

keeps pH within normal limits Resist pH change by taking up excess H+ or

OH-

pH of blood is about 7.4 – maintained by buffer

Figure 2.17 The pH scaleCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

Acid

Base

Increasin

g [H

+]In

creasing

[OH

–]

0

1

2

3

4

5

6

7

8

9

10

11

12

13

14

hydrochloric acid (HCI)

stomach acid

lemon juice

Coca-Cola, beer, vinegar

tomatoes

black coffee

urine

pure water, tears

seawater

baking soda, stomach antacids

Great Salt Lake

household ammonia

bicarbonate of soda

oven cleaner

sodium hydroxide (NaOH)

normal rainwater

saliva

human blood

milk of magnesia

[H+]

[OH–]

neutral pH[H+] =

[OH–]

Self-test/Review Questions

Use these questions as a self test and then discuss your responses with your study group/classmates—your responses will not be collected.

1. Why is carbon dioxide gas, CO2, classified as a compound but nitrogen gas, N2, is not?

2. Which of the following are compounds? Elements?: C6H12O6, CH4, O2, Cl2, HCl, MgCl2, Fe, Ca, Ne, NaI, I

3. What is the difference between an atom and an ion? Give examples of each to support your response.

4. Which subatomic particle determines the identity of an atom?

5. Which subatomic particle determines the chemical properties of an atom?

Self-test/Review Questions

6. A carbon atom has 6 protons, and the most common isotope of carbon has 6 neutrons. A radioactive isotope of carbon has 8 neutrons. What are the atomic numbers and the mass numbers of the of the stable and radioactive forms of carbon?

7. Explain the difference between an ionic and covalent bond in terms of what happens to the electrons in the outer shell of the participating atoms.

8. Sodium fluoride, NaF, is often added to toothpaste to both kill bacteria that cause cavities. It also helps to harden the enamel of teeth thus helping it resist cavities. Is sodium fluoride an ionic or covalent compound? How do you know? Explain your reasoning.

9. Is carbon dioxide an ionic or covalent compound? How do you know? Explain your reasoning.

Self-test/Review Questions (cont.)

10. Why are the following incorrect structures for the substances below? Rewrite their structures with the correct number of chemical bonds.

a. Carbon dioxide gas: O—C—O

b. Oxygen gas: O—O

c. Nitrogen gas: N—N

11. Explain how water’s versatility as a solvent results from the fact that water is polar molecule.

12. A bottle of Pepsi consists mostly of sugar dissolved in water, with some carbon dioxide gas that makes fizzy and makes the pH less than 7. Describe Pepsi using the following terms: solute, solvent, acidic, aqueous solution


13. Which of the following are chemical changes? Physical changes? If possible, write the balanced chemical equation for those that are a chemical change.

a. The alcoholic fermentation in Yeast in which yeast produce ethanol, C2H5OH, and carbon dioxide, CO2, from the sugar glucose, C6H12O6

b. Water boils to form steam

c. The healing of a cut finger

d. Cutting a piece of wood with a saw

e. Potassium metal, K, and chlorine gas (Cl2) combine to form potassium chloride.

f. The rusting of iron, Fe, to produce rust, iron (III) oxide (Fe2O3)


14. Which of these is not a subatomic particle? a) proton; b) ion; c) neutron; d) electron

15. The outermost electron shell of every Noble Gas element (except Helium) has ___ electrons. a) 1; b) 2; c) 4; d) 6; e) 8

16. An organic molecule is likely to contain all of these elements except ___. a) C; b) H; c) O; d) Ne; e) N

17. The chemical bond between water molecules is a ___ bond. a) ionic; b) polar covalent; c) nonpolar covalent; d) hydrogen

18. A solution with a pH of 7 has ___ times more H ions than a solution of pH 9. a) 2; b) 100; c) 1000; d) 9; e) 90

19. The type of chemical bond formed when electrons are shared between atoms is a ___ bond. a) ionic; b) covalent; c) hydrogen


20. The type of chemical bond formed when oppositely charged particles are attached to each other is a ___ bond. a) ionic; b) covalent; c) hydrogen

21. Carbon has an atomic number of 6. This means it has ___. a) six protons; b) six neutrons; c) six protons plus six neutrons; d) six neutrons and six electrons

22. Each of the isotopes of hydrogen has ___ proton(s). a) 3; b) 1; c) 2; d) 92; e) 1/2

23. A molecule is ___. a) a mixture of various components that can vary; b) a combination of many atoms that will have different ratios; c) a combination of one or more atoms that will have a fixed ratio of its components; d) more important in a chemistry class than in a biology class

Estimating the Size of an Object Viewed with a Microscope

• Calculate the length and width of the following microscopic object in both millimeters (mm) and micrometers (m). 1 mm = 1000 m

• Base your calculations on the following field sizes:

Low power (40x): 4.5 mm

Medium power (100x): 1.8 mm

High power (400x): 0.45 mm

Object viewed at medium power (100x)

Remember: Field size decreases by the same factor as the magnification increases!

Estimating the Size of an Object Viewed with a Microscope

• Calculate the length and width of the following microscopic object in both millimeters and micrometers. 1 mm = 1000 m

• Base your calculations on the following hypothetical field sizes:

Low power (30x): 4.0 mm = ___m

Medium power (180x): ___mm = ___m

High power (300x): ___mm = ___m

Object viewed at high power (300x)

Remember: Field size decreases by the same factor as the magnification increases!

Essentials of Biology Sylvia S. Mader Chapter 2 Lecture Outline Copyright © The McGraw-Hill...

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Transcript of Essentials of Biology Sylvia S. Mader Chapter 2 Lecture Outline Copyright © The McGraw-Hill...