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Chapter 2The Chemical Context of Life

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Matter

• Takes up space and has mass

• Exists as elements (pure form) and in chemical combinations called compounds

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Elements

• Can’t be broken down into simpler substances by chemical reaction

• Composed of atoms• Essential elements in living things

include carbon C, hydrogen H, oxygen O, and nitrogen N making up 96% of an organism

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Other Elements• A few other elements Make up the

remaining 4% of living matter

Table 2.1

Trace Elements

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Deficiencies

(a) Nitrogen deficiency(b) Iodine deficiency (Goiter)

• If there is a deficiency of an essential element, disease results

Figure 2.3

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Trace Elements

• Trace elements Are required by an organism in only minute quantities

• Minerals such as Fe and Zn are trace elements

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Compounds

Sodium Chloride Sodium Chloride

+

• Are substances consisting of two or more elements combined in a fixed ratio

• Have characteristics different from those of their elements

Figure 2.2

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Properties of Matter• An element’s

properties depend on the structure of its atoms

• Each element consists of a certain kind of atom that is different from those of other elements

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

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Subatomic Particles

• Atoms of each element Are composed of even smaller parts called subatomic particles

• Neutrons, which have no electrical charge

• Protons, which are positively charged

• Electrons, which are negatively charged

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Subatomic Particle Location

• Protons and neutrons– Are found in the

atomic nucleus• Electrons

– Surround the nucleus in a “cloud”

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Simplified models of an Atom

Nucleus

(a) (b)In this even more simplifiedmodel, the electrons areshown as two small bluespheres on a circle around thenucleus.

Cloud of negativecharge (2 electrons)

Electrons

This model represents theelectrons as a cloud ofnegative charge, as if we hadtaken many snapshots of the 2electrons over time, with eachdot representing an electron‘sposition at one point in time.

Figure 2.4

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Atomic Number & Atomic Mass

• Atoms of the various elements Differ in their number of subatomic particles

• The number of protons in the nucleus = atomic number

• The number of protons + neutrons = atomic mass

• Neutral atoms have equal numbers of protons & electrons (+ and – charges)

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Atomic Number

•Is unique to each element and is used to arrange atoms on the Periodic table•Carbon = 12•Oxygen = 16•Hydrogen = 1•Nitrogen = 17

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Atomic Mass

• Is an approximation of the atomic mass of an atom•It is the average of the mass of all isotopes of that particular element•Can be used to find the number of neutrons (Subtract atomic number from atomic mass)

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Isotopes• Different forms of the same element• Have the same number of protons, but

different number of neutrons• May be radioactive spontaneously

giving off particles and energy• May be used to date fossils or as

medical tracers

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Energy Levels of Electrons• An atom’s electrons Vary in the

amount of energy they possess• Electrons further from the nucleus

have more energy• Electron’s can absorb energy and

become “excited” • Excited electrons gain energy and

move to higher energy levels or lose energy and move to lower levels

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Energy

• Energy– Is defined as the capacity to cause

change• Potential energy - Is the energy that matter possesses

because of its location or structure• Kinetic Energy - Is the energy of motion

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Electrons and Energy

• The electrons of an atom– Differ in the amounts of potential energy

they possess

A ball bouncing down a flightof stairs provides an analogyfor energy levels of electrons,because the ball can only reston each step, not betweensteps.

(a)

Figure 2.7A

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Energy Levels• Are represented by electron shells

Third energy level (shell)

Second energy level (shell)

First energy level (shell)

Energyabsorbed

Energylost

An electron can move from one level to another only if the energyit gains or loses is exactly equal to the difference in energy betweenthe two levels. Arrows indicate some of the step-wise changes inpotential energy that are possible.

(b)

Atomic nucleus

Figure 2.7B

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Thermodynamics and BiologyFirst law of thermodynamics• In any process, the total energy of the universe remains the

same.• It can also be defined as: for a thermodynamic cycle the

sum of net heat supplied to the system and the net work done by the system is equal to zero.

Second law of thermodynamics• The entropy (useless energy) of an isolated system will tend

to increase over time.• In a simple manner, the second law states that "energy

systems have a tendency to increase their entropy" rather than decrease it.

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Periodic table– Shows the electron distribution for

all the elements

Secondshell

Helium2He

Firstshell

Thirdshell

Hydrogen1H

2He

4.00Atomic mass

Atomic number

Element symbol

Electron-shelldiagram

Lithium3Li

Beryllium4Be

Boron3B

Carbon6C

Nitrogen7N

Oxygen8O

Fluorine9F

Neon10Ne

Sodium11Na

Magnesium12Mg

Aluminum13Al

Silicon14Si

Phosphorus15P

Sulfur16S

Chlorine17Cl

Argon18Ar

Figure 2.8

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Why do some elements react?• Valence electrons

– Are those in the outermost, or valence shell

– Determine the chemical behavior of an atom

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Electron Orbitals• An orbital

– Is the three-dimensional space where an electron is found 90% of the time

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Covalent Bonds

Figure 2.10

• Sharing of a pair of valence electrons

• Examples: H2

Hydrogen atoms (2 H)

Hydrogenmolecule (H2)

+ +

+ +

+ +

In each hydrogenatom, the single electronis held in its orbital byits attraction to theproton in the nucleus.

1

When two hydrogenatoms approach eachother, the electron ofeach atom is alsoattracted to the protonin the other nucleus.

2

The two electronsbecome shared in a covalent bond,forming an H2

molecule.

3

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Covalent Bonding• A molecule

– Consists of two or more atoms held together by covalent bonds

• A single bond– Is the sharing of one pair of valence electrons

• A double bond– Is the sharing of two pairs of valence electrons

Name(molecularformula)

Electron-

shelldiagram

Structuralformula

Space-fillingmodel

Hydrogen (H2). Two hydrogen atoms can form a single bond.

Oxygen (O2). Two oxygen atoms share two pairs of electrons to form a double bond.

H H

O O

Figure 2.11 A, B

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Compounds & Covalent Bonds

Name(molecularformula)

Electron-shell

diagram

Structuralformula

Space-fillingmodel

(c)

Methane (CH4). Four hydrogen atoms can satisfy the valence ofone carbonatom, formingmethane.

Water (H2O). Two hydrogenatoms and one oxygen atom arejoined by covalent bonds to produce a molecule of water.

(d)

HO

H

H H

H

H

C

Figure 2.11 C, D

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Covalent Bonding

• In a nonpolar covalent bond– The atoms have

similar electronegativities

– Share the electron equally

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Figure 2.12

This results in a partial negative charge on theoxygen and apartial positivecharge onthe hydrogens.

H2O

d–

O

H Hd+ d+

Because oxygen (O) is more electronegative than hydrogen (H), shared electrons are pulled more toward oxygen.

• In a polar covalent bond– The atoms have differing

electronegativities– Share the electrons unequally

Covalent Bonding

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Ionic Bonds

• In some cases, atoms strip electrons away from their bonding partners

• Electron transfer between two atoms creates ions

• Ions– Are atoms with more or fewer electrons

than usual– Are charged atoms

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Ions

• An anion– Is negatively

charged ions• A cation

– Is positively charged

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Ionic Bonding

Cl–

Chloride ion(an anion)

–

The lone valence electron of a sodiumatom is transferred to join the 7 valenceelectrons of a chlorine atom.

1 Each resulting ion has a completedvalence shell. An ionic bond can formbetween the oppositely charged ions.

2

Na NaCl Cl

+

NaSodium atom(an uncharged

atom)

ClChlorine atom(an uncharged

atom)

Na+

Sodium on(a cation)

Sodium chloride (NaCl)

Figure 2.13

• An ionic bond– Is an attraction between anions and

cations

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Ionic Substances

Na+

Cl–Figure 2.14

• Ionic compounds– Are often

called salts, which may form crystals

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Weak Chemical Bonds

• Several types of weak chemical bonds are important in living systems

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Hydrogen Bonds

– +

+

Water(H2O)

Ammonia(NH3)

OH

H +

–

N

HH H

A hydrogenbond results from the attraction between thepartial positive charge on the hydrogen atom of water and the partial negative charge on the nitrogen atom of ammonia.+ d+

Figure 2.15

• A hydrogen bond– Forms when a hydrogen atom covalently

bonded to one electronegative atom is also attracted to another electronegative atom

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Van der Waals Interactions• Van der Waals interactions

– Occur when transiently positive and negative regions of molecules attract each other

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Weak Bonds

• Weak chemical bonds– Reinforce the shapes of large molecules– Help molecules adhere to each other

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Molecular Shape and Function

• Structure determines Function!• The precise shape of a molecule

– Is usually very important to its function in the living cell

– Is determined by the positions of its atoms’ valence orbitals

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Orbitals & Covalent Bonds

Space-fillingmodel

Hybrid-orbital model(with ball-and-stick

model superimposed)UnbondedElectron pair

104.5°

O

HWater (H2O)

Methane (CH4)

H

H H

H

C

O

H

H

H

C

Ball-and-stickmodel

H H

H

H

(b) Molecular shape models. Three models representing molecular shape are shown for two examples; water and methane. The positions of the hybrid orbital determine the shapes of the moleculesFigure 2.16 (b)

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Shape and Function

• Molecular shape– Determines how biological molecules

recognize and respond to one another with specificity

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Morphine

Carbon

Hydrogen

Nitrogen

Sulfur

OxygenNaturalendorphin

(a) Structures of endorphin and morphine. The boxed portion of the endorphin molecule (left) binds toreceptor molecules on target cells in the brain. The boxed portion of the morphine molecule is a close match.

(b) Binding to endorphin receptors. Endorphin receptors on the surface of a brain cell recognize and can bind to both endorphin and morphine.

Naturalendorphin

Endorphinreceptors

Morphine

Brain cell

Figure 2.17

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Chemical Reactions

• Chemical reactions make and break chemical bonds

• A Chemical reaction– Is the making and breaking of chemical

bonds– Leads to changes in the composition of

matter

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Chemical Reactions

Reactants Reaction Product

2 H2 O2

2 H2O

+

+

• Chemical reactions– Convert reactants to products

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Chemical Reactions• Photosynthesis

– Is an example of a chemical reaction

Figure 2.18

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