Essential Chemistry for Biology - Arizona State...

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Copyright © 2004 Pearson Education, Inc. publishing as Benjamin Cummings

PowerPoint® Lecture Slides forEssential Biology, Second Edition & Essential Biology with Physiology

Neil Campbell, Jane Reece, and Eric Simon

Presentation prepared by Chris C. Romero

CHAPTER 2CHAPTER 2

Essential Chemistry for BiologyFigures 2.1 – 2.7


• Because of surface tension, some insects can walk on water. Water molecules “stick” to each other by cohesionwhich results from hydrogen bonding.

• The iron in a multivitamin pill is the same element as the iron in a train or ship


• There has been a sharp decline in tooth decay in the last few decades

BIOLOGY AND SOCIETY:

FLUORIDE IN THE WATER

– Fluoride-containing chemicals have been added to drinking water and dental products

• The use of fluoride in drinking water illustrates the point that organisms are chemical systems

2


• Biology includes the study of life at many levels

TRACING LIFE DOWN TO THE CHEMICAL LEVEL

• In order to understand life, we will start at the macroscopic level, the ecosystem, and work our way down to the microscopic level of cells

• Cells consist of enormous numbers of chemicals that give the cell the properties we recognize as life


Figure 2.1

Ecosystem African savanna

CommunityAll organisms in savanna

PopulationHerd of zebrasOrganism Zebra

Organ systemCirculatory system

OrganHeart

CellHeart muscle cell

TissueHeart muscletissue

MoleculeDNA

AtomOxygen atom


Ecosystem

Community

Population ex. all humans in city, all termites in class

Individual Organism

Organ Systems ex. respiratory, reproductive, circulatory

Organs ex. lungs, ovaries, heart

Tissue ex. connective, nervous, muscular

Cells ex. neuron, sarcomere, epithelial

Organelles ex, nucleus, chloroplast, mitochondria

Macromolecules ex. DNA, RNA, cellulose, lipids

3


• Take any biological system apart and you eventually end up at the chemical level.

SOME BASIC CHEMISTRY

Cells ex. Prokaryotic, Eukaryotic

Macromolecules ex. DNA, RNA, fat

Molecules ex. H2O, HCl, H2SO4,

Atoms ex. C, H, O, N, Iodine C=carbon

Subatomic particles: within nucleus (neutron & proton)around nucleus (electrons)


• Matter is anything that occupies space and has mass

Matter: Elements and Compounds

• Matter is found on the Earth in “3” physical states.

– Solid

– Liquid

– Gas

– Plasma


• Matter is composed of chemical elements.

– Elements are substances that cannot be broken down into other substances

– There are 92 naturally occurring elements on Earth

4


• All the elements are listed in the periodic table.

Atomic number

Element symbol

Mass number

Figure 2.2


• Twenty-five elements are essential to life.

– Four of these make up about 96% of the weight of the human body C,H,O,N

– Trace elements occur in smaller amounts

Figure 2.3


• Trace elements are essential for life

– An iodine deficiency causes goiter.

Figure 2.4

5


• Elements can combine to form compounds

– These are substances that contain two or more elements in a fixed ratio

– Example: NaCl (salt).


• Each element consists of one kind of atom

Atoms

– An atom is the smallest unit of matter that still retains the properties of an element.

Nucleus

Cloud of negativecharge (2 electrons)

(a)

(b)

2 Protons

2 Neutrons

2 Electrons Figure 2.5


• Atoms are composed of subatomic particles

The Structure of Atoms

– A proton is positively charged

– An electron is negatively charged

– A neutron is electrically neutral.

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• Most atoms have protons and neutrons packed tightly into the nucleus

– The nucleus is the atom’s central core

– The electrons orbit the nucleus


• Elements differ in the number of subatomic particles in their atoms

– The number of protons, the atomic number, determines which element it is

– An atom’s mass number is the sum of the number of protons and neutrons

– Mass is a measure of the amount of matter in an object; protons and neutrons each have an atomic mass unit of 1


• Isotopes are alternate mass forms of an element

Isotopes

– They have the same number of protons and electrons

– But they have a different number of neutrons, p.23

Table 2.1, p. 23

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• Radioactive isotopes

– The nucleus decays, giving off particles and energy

• Radioactive isotopes have many uses in research and medicine

– Example: PET scans


Figure 2.6

Hearingwords

Seeingwords

Speakingwords

Generatingwords

(b)(a)


• Uncontrolled exposure to radioactive isotopes can harm living organisms by damaging DNA

– Example: the 1999 Tokaimura nuclear accident

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• Electrons determine how an atom behaves when it encounters other atoms

Electron Arrangement and the Chemical Properties of Atoms

• Electrons orbit the nucleus of an atom in specific electron shells

– The number of electrons in the outermost shell determines the chemical properties of an atom.


• Atoms of the four elements most abundant in life.

Figure 2.7

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


• Chemical reactions enable atoms to give up or acquire electrons in order to complete their outer shells

Chemical Bonding and Molecules

– These interactions usually result in atoms staying close together

– The atoms are held together by chemical bonds.

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

Ionic Bonds

– Charged atoms are called ions

– Ionic bondsare formed between oppositely charged ions.

Figure 2.8

Sodium atom (Na) Chlorine atom (Cl)

Completeouter shells

Sodium ion (Na ++++) Chloride ion (Cl −−−−)

Sodium chloride (NaCl)


• A covalentbond forms when two atoms share one or more pairs of outer-shell electrons.ex.withinwater molecule

Covalent Bonds

Figure 2.9


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

Chemical Reactions

– Such changes in the chemical composition of matter are called chemical reactions How many oxygen atoms as products?

Unnumbered Figure 2.1

Hydrogen gas Oxygen gas Water

Reactants Products

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• Chemical reactions can be symbolized with equations

– “On the left side of the equation are the reactants, the starting materials”

– “On the right side of the equation are the products, the end materials”

The arrow always points to the products!


• Chemical reactions cannot create or destroy matter

– They only rearrange it

(Law of Conservation of Energy)


• According to the text life on Earth began in water and evolved there for 3 billion years

WATER AND LIFE

– Modern life still remains tied to water

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

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

Figure 2.10


• Studied in isolation, the water molecule is deceptively simple

The Structure of Water

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

Unnumbered Figure 2.2

H

O

H


• But the electrons of the covalent bondsare not shared equally between oxygen and hydrogen

– This unequal sharing makes water a polar molecule

(++++) (++++)

(−−−−) (−−−−)

Figure 2.11a

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• The polarity of water results in weak electrical attractions between neighboring water molecules. These interactions are called hydrogen bondsand result in cohesionwhich accounts for surface tension

(b)

(−−−−)

Hydrogen bond(++++)

(++++)(−−−−)

(−−−−)

(++++)

(++++)(−−−−)

Figure 2.11b


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

Water’s Life-Supporting Properties

– Water’s cohesivenature

– Water’s ability to moderate temperature

– Floating ice D=M/V, see p. 30

– Versatility of water as a solvent.


• Water molecules stick together as a result of hydrogen bonding

The Cohesion of Water

– This is called cohesion

– Cohesion is vital for water transport in plants.

Figure 2.12

Microscopic tubes

13


• Surface tensionis 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.

How Water Moderates Temperature


• Heat and temperature are related, but different

– Heat is the amount of energy associated with the movement of the atoms and molecules in a body of matter

– Temperature measures the intensity of heat

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

14


• Water can moderate temperatures

– Earth’s giant water supply causes temperatures to stay within limits that permit life

– Evaporative cooling removes heat from the Earth and from organisms

Figure 2.14


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

The Biological Significance of Ice Floating

– A chunk of ice has fewer molecules than an equal volume of liquid water, p. 30


• The density of ice is lower than liquid water

– This is why ice floats

Figure 2.15

Hydrogen bond

Liquid water

Hydrogen bondsconstantly break and re-form

Ice

Stable hydrogen bonds

15


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

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


• 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, p. 30

– The dissolved substance is called the solute

Figure 2.16

Ion in solutionSalt crystal


• When water is the solvent, the result is called an aqueous solution. Water is a very common solvent.

16


• Acid

Acids, Bases, and pH

– A chemical compound that donates H+ ions to solutions. Acids are strong if pH near 1 and weak if pH near to 7. ex. HCl, H2SO4

• Base

– A compound that accepts H+ ions and removes them from solution. Strong bases have pH near 14, weak ones near 7.


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

pH scale

• To describe the acidity of a solution, we use the pH scale

Figure 2.17


• Buffers are substances that resist pH change

– They accept H+ ions when they are in excess

– They donate H+ ions when they are depleted

• Buffering is not foolproof

– Example: acid precipitation.

Figure 2.18

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• Chemical reactions and physical processes on the early Earth created an environment that made life possible.

EARTH BEFORE LIFE


• It is estimated that the Earth began as a cold world about 4.5 billion years ago

• The planet eventually melted from heat produced by

– compaction

– radioactive decay

– impact of meteorites.


• The first atmosphere was probably composed of hot hydrogen gas

– This gas escaped because the gravity of Earth was not strong enough

• A new atmosphere was formed from the gases belched from volcanoes.

Figure 2.19

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• Atoms

SUMMARY OF KEY CONCEPTS

Visual Summary 2.1

Proton

•••• Positive charge

•••• Determines element

Neutron

•••• No charge

•••• Determines isotope

Electron

•••• Negative charge

• Participates in chemicalreactions

• Outer-shell electronsdetermine chemicalbehavior

Nucleus

• Consists of neutrons and protons

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