Chapter 2 ppt

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Biology

Sylvia S. Mader Michael Windelspecht

Chapter 2 Basic Chemistry

Lecture Outline

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See separate FlexArt PowerPoint slides for all figures and tables pre-inserted into

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2.1 Chemical Elements

• Matter is defined as anything that has mass

and occupies space.

• Matter exists in three states: solid, liquid, and gas.

• All matter (both living and non-living) is

composed of elements.

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Elements

• An element is a substance that cannot be

broken down into substances with different properties; composed of one type of atom.

• Six elements make up 95% of the body weight of organisms: (acronym CHNOPS)

 Carbon

 Hydrogen

 Nitrogen

 Oxygen

 Phosphorus

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4

60

40

20

0

Fe Ca K S P Si Al Mg Na O N C H

Earth’s crust organisms

Element

P

e

rc

e

nt

by

W

e

ight

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Atoms

• An atom is the smallest part of an element

that displays the property of the element.

• An element and its atom share the same name.

 Composed of subatomic particles: protons, neutrons, electrons

 Central nucleus

• Protons- positively charged, 1 amu

• Neutrons- no charge, 1 amu

 Orbiting clouds around nucleus (electron shells)

• Electrons- negatively charged, very low

mass-negligible in calculations

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c.

Subatomic Particles

= proton

b. a.

Particle

Proton +1 1 Nucleus Atomic Mass Unit

(AMU) Location Electric

Charge

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c.

Subatomic Particles

= proton = neutron

b. a.

Particle Proton Neutron

Nucleus Nucleus +1

0

1 1

Atomic Mass Unit

Charge

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8 c. Subatomic Particles = proton = neutron = electron b. a. Particle Proton Neutron Electron Nucleus Nucleus Electron shell +1 0 –1 1 1 0

Atomic Mass Unit

Charge

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Atomic Number and Mass Number

• Each element is represented by one or two

letters to give it a unique atomic symbol.

 H = hydrogen, Na = sodium, C = carbon

• The atomic number is equal to the number of protons in each atom of an element.

• The mass number of an atom is equal to the

sum of the number of protons and neutrons in atom’s nucleus.

 The atomic mass is approximately equal to the mass

number.

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mass number

atomic number

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Periodic Table

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• Atoms of an element are arranged horizontally by increasing atomic number in rows called

periods.

• Atoms of an element arranged in vertical columns are called groups.

 Atoms within the same group share the same binding characteristics.

• Atoms shown in the periodic table are electrically neutral.

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Pe riod s 1 2 3 4 atomic number

atomic symbol atomic mass

VIII Groups 20.18 19.00 16.00 14.01 12.01 10.81 9.012 6.941

22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95

83.60 79.90 78.96 74.92 72.59 69.72 40.08 39.10 Kr Br Se As Ge Ga Ca K 36

19 20 31 32 33 34 35

11 12 13 14 15 16 17 18

Na Mg Al Si P S Cl Ar

Ne F O N C B Be Li 10 9 8 7 6 5 4 3 4.003 He 2 VII VI V IV III II 1.008 I 1 H

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Isotopes

• Isotopes are atoms of the same element that differ in the number of neutrons (and therefore different atomic masses).

 Some isotopes spontaneously decay

• Radioactive isotopes give off energy in the form of rays and subatomic particles

• Can be helpful or harmful

12 6

Carbon 12

C 13 ₆

Carbon 13

C 14 ₆

Carbon 14

C

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b. a.

larynx

thyroid gland trachea

a: © Biomed Commun./Custom Medical Stock Photo; b (left) : Courtesy National Institutes of Health; b: (right) © Mazzlota et al./Photo Researchers, Inc

Uses of Low

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Uses of High Levels of Radiation

a: (Peaches): © Tony Freeman/PhotoEdit; b: © Geoff Tompkinson/SPL/Photo Researchers, Inc.

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Electrons and Energy

• Electrons are attracted to the positively charged nucleus, thus it takes energy to hold electrons in place.

• It takes energy to push them away and keep them in their own shell.

 The more distant the shell, the more energy it takes to hold in place.

• Electrons have energy due to their relative position (potential energy).

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The Distribution of Electrons

• The Bohr model is a useful way to visualize electron location.

 Electrons revolve around the nucleus in energy shells (energy levels).

 For atoms with atomic numbers of 20 or less, the following rules apply:

• the first energy shell can hold up to 2 electrons

• each additional shell can hold up to 8 electrons

• each lower shell is filled first before electrons are placed in the next shell

 These rules cover most of the biologically

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Valence Electrons

• The outermost energy shell is called the

valence shell.

• The valence shell is important because it determines many of an atom’s chemical properties.

• The octet rule states that the outermost shell is

most stable when it has eight electrons.

 Exception: If an atom has only one shell, the

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Valence Electrons

• The number of electrons in an atom’s valence

shell determines whether the atom gives up,

accepts, or shares electrons to acquire eight electrons in the outer shell.

–Atoms that have their valence shells filled with electrons tend to be

chemically stable.

–Atoms that do not have their valence

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H

P S

C

O

N

electron valence shell

nitrogen carbon hydrogen sulfur phosphorus 32 16 15 31 8 16 14 7 6 12 1

1 H C N

S P O oxygen electron shell nucleus

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2.2 Compounds and Molecules

• A molecule is two or more elements bonded

together.

• A compound is a molecule containing at least two different elements bonded together.

 CO₂, H₂O, C₆H₁₂O₆, etc.

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Compounds and Molecules

one molecule

indicates 6 atoms of carbon

indicates 6 atoms of oxygen

indicates 12 atoms of hydrogen

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Chemical Bonding

• Bonds that exist between atoms in

molecules contain energy.

• Bonds between atoms are caused by the

interactions between electrons in

outermost energy shells.

• The process of bond formation is called a

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Types of Bonds: Ionic Bonding

• An ion is an atom that has lost or gained an

electron.

• An ionic bond forms when electrons are

transferred from one atom to another atom and the oppositely charged ions are attracted to one.

 Example: formation of sodium chloride

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Na Cl

sodium atom (Na) chlorine atom (Cl)

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Na Cl

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+ –

Na Cl

sodium ion (Na+) chloride ion (Cl–)

sodium chloride (NaCl)

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+ –

Na

sodium ion (Na+₎

sodium chloride (NaCl)

a. b.

chloride ion (Cl–)

Na Cl

Cl– Na+

Cl

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Types of Bonds: Covalent Bonds

• Covalent bonds result when two atoms share electrons so each atom has an octet of electrons in the outer shell.

 Note: In the case of hydrogen, the outer energy shell is complete when it contains 2 electrons.

• In a nonpolar covalent bond electrons are

shared equally between atoms.

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30 a. Hydrogen gas

b. Oxygen gas

c. Methane

Structural Formula Electron Model

H C H H H H H O O C O O H H H H H H CH₄ O₂ H₂ Molecular Formula

Covalently

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Types of Bonds: Covalent Bonds

• In a polar covalent bond electrons are shared

unequally.

 Example: water

• Electronegativity is the ability of an atom to attract electrons towards itself in a chemical bond.

 In water, the oxygen atom is more electronegative than the hydrogen atoms and the bonds are polar.

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O

H

Oxygen is partially negative ( )

Hydrogens are partially positive ( )

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2.3 Chemistry of Water

• Water is a polar molecule.

 The shape of a water molecule and its polarity make hydrogen bonding possible.

• A hydrogen bond is a weak attraction between a slightly positive hydrogen atom and a slightly negative atom.

 Can occur between atoms of different molecules or within the same molecule

 A single hydrogen bond is easily broken while

multiple hydrogen bonds are collectively quite strong.

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Fig. 2.9

H

H H H

H H H H

O O O O Electron Model Ball-and-stick Model Space-filling Model

Oxygen attracts the shared electrons and is partially negative.

b. Hydrogen bonding between water molecules hydrogen

bond

Hydrogens are partially positive.

a. Water (H₂O)

104.5˚ + + - + + -

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Properties of Water

• Water molecules cling together because of hydrogen bonding.

 This association gives water many of its unique chemical properties.

• Water has a high heat capacity.

 The presence of many hydrogen bonds allow water to absorb a large amount of thermal heat without a great change in temperature.

 The temperature of water rises and falls slowly. • Allows organisms to maintain internal

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Cal orie s of Hea t En e rgy / g 0 800 600 400 200 Solid evaporation occurs Liquid 540 calories Gas 80 calories

a. Calories lost when 1 g of liquid water freezes and calories required when 1 g of liquid water evaporates.

b. Bodies of organisms cool when their heat is used to evaporate water. freezing occurs

Temperature (°C)

120 100 80 0 20 40 60

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Properties of Water

• Water has a high heat of vaporization.

 Hydrogen bonds must be broken to evaporate water.

 Bodies of organisms cool when their heat is used to evaporate water.

• Water is a good solvent.

 Water is a good solvent because of its polarity.

 Polar substances dissolve readily in water.

 Hydrophilic molecules dissolve in water.

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H

H H H H H

H H H

H

An ionic salt

dissolves in water.

H H

Cl– Na+

O

O O

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Properties of Water

• Water molecules are cohesive and adhesive.

 Cohesion is the ability of water molecules to cling to each other due to hydrogen bonding.

• Water flows freely • Surface tension

 Adhesion is the ability of water molecules to cling to other polar surfaces.

• Adhesion is due to water’s polarity.

• Capillary action

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Water as a Transport Medium

Water evaporates, pulling the water column from the roots to the leaves.

Water molecules cling together and adhere to sides of vessels in stems.

Water enters a plant at root cells.

H2O

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Properties of Water

• Ice is less dense than liquid water.

 At temperatures below 4°C, hydrogen bonds

between water molecules become more rigid but also more open.

 Water expands as it reaches 0°C and freezes.

 Ice floats on liquid water.

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Density of Water at Various

Temperatures

0 4 100

1.0

0.9

Den

s

it

y

(g

/c

m

3 )

Temperature (ºC)

liquid water ice lattice

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A Pond in Winter

ice layer

Protists provide food for fish.

River otters visit ice-covered ponds.

Aquatic insects survive in air pockets.

Freshwater fish take oxygen from water.

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2.4 Acids and Bases

• pH is a measure of hydrogen ion concentration

in a solution.

• When water ionizes or dissociates, it releases an

equal number of hydrogen (H+_{) ions and}

hydroxide (OH-) ions.

• Acids are substances that dissociate in water, releasing hydrogen ions.

• Bases are substances that either take up

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The pH Scale

• The pH scale is used to indicate the acidity or

basicity (alkalinity) of a solution.

 Values range from 0-14 • 0 to <7 = Acidic

• 7 = Neutral

• >7 to 14 = Basic (or alkaline)

 Logarithmic scale

• Each unit change in pH represents a 10-fold change in H+ concentration

• pH of 4 is 10X as acidic as pH of 5

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1 0 2

3

4 5

6 7 8 ₉

13 10 11 12 14 pu re w ater , tear s

Acidic _Basic

OH–

N

eutral

pH

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Buffers and pH

• A buffer is a chemical or a combination of

chemicals that keeps pH within normal limits.

• Health of organisms requires maintaining the pH of body fluids within narrow limits.

 Human blood is normally pH 7.4 (slightly basic) • If blood pH drops below 7.0, acidosis results • If blood pH rises above 7.8, alkalosis results • Both are life-threatening situations.