C H A P T E R 2

THE CHEMISTRY OF LIFE

ATOMS

• Any substance that has mass and occupies space is called matter.

• Matter is composed of small particles called atoms.

• An atom is the smallest particle that still retains the chemical properties of a substance.

MASS VS. WEIGHT

• Mass – the amount of a substance.• Weight – the force gravity exerts on a

substance.• An objects mass is the same on Earth or

the moon.• An objects weight will be greater on Earth

because the Earth’s gravitational force is greater.

BASIC ATOMIC STRUCTURE

• Core nucleus containing protons (+) and neutrons.

• Cloud of electrons (-) surrounding the nucleus.• Protons, neutrons, and electrons are called

subatomic particles.

IONS

• Neutral atom (no charge) one orbiting electron (-) for every proton (+) in the nucleus.

• Atoms where the protons and electrons do not balance are called ions.

• Ions may be positively or negatively charged.

ELECTRONS

• Protons and neutrons contribute most of the mass of an atom – electrons have very little mass.

• Electrons determine chemical behavior of atoms because they come into contact with electrons of other atoms.

ELECTRONS CARRY ENERGY

• It takes work to hold electrons in orbit.• Electrons have energy of position – potential energy.

ELECTRONS CARRY ENERGY

• An electron can be transferred from one atom to another.

• Loss of an electron is called oxidation.• Gain of an electron is called reduction.

ELECTRONS CARRY ENERGY

• We use simplistic models to visualize atoms.

• The volume of space around a nucleus where an electron is most likely to be found is called the orbital of that electron.

ELECTRONS CARRY ENERGY

• Each energy level has a specific number of orbitals, which can hold up to 2 electrons.• The first energy level has 1 orbital.• The outer energy levels have 4 orbitals each.

• Atoms with incomplete orbitals tend to be more reactive.

ATOMIC NUMBER

• The number of protons in the nucleus is the atomic number.• Atomic number of sodium is 11.

ATOMIC MASS

• Neutrons and protons are similar in mass.• Each weighs about 1 dalton.• The number of neutrons plus the number of

protons is called the atomic mass.• Atomic mass of sodium is about 22 (22.989)

daltons.

• Electrons have negligible mass.

ELEMENTS

• Atoms with the same atomic number (# of protons) have the same chemical properties and belong to the same element.

ISOTOPES

• Atoms that have the same number of protons, but different number of neutrons are called isotopes.• Same atomic number, different mass.

ISOTOPES

• Example – Carbon: 6 protons, 6 electrons.• Usually has 6 neutrons = Carbon-12• Sometimes has 7 or 8 neutrons.• Carbon-14 is unstable; nucleus tends to break

up (radioactive decay).

DATING FOSSILS

• Fossils are created when remains, footprints, or other traces become buried in sand or sediment.

• Over time, calcium is mineralized as the sediment turns to rock.

• A fossil is any record of prehistoric life – generally over 10,000 years old.

DATING FOSSILS

• Scientists date the rocks containing fossils to get an idea of how old the fossils are.

• Rocks can be dated by measuring the degree of radioactive decay of certain radioactive atoms among rock-forming minerals.

DATING FOSSILS

• A radioactive element contains so many neutrons and protons that it is unstable and likely to fly apart.

• The rate of decay for a particular element is constant and so it can be used to date rocks.

DATING FOSSILS

• Carbon-14 (14C) radioisotopic dating is often used to date fossils less than 50,000 years old.

• A certain proportion of naturally occurring carbon is carbon-14 rather than carbon-12.

• After an organism dies, its 14C decays over time.

DATING FOSSILS

• It takes 5,600 years for half of the 14C to be converted to 14N.

• This is called the half-life of the isotope.• Constant

• A sample with ¼ the original proportion of 14C would be 11,200 years old (2 half-lives).

• 40K decay into 40Ar has a half-life of 1.3 billion years.

MOLECULES

• A molecule is a group of atoms held together by energy.

• The force holding two atoms together is called a chemical bond.• Ionic bond• Covalent bond• Hydrogen bond

IONIC BONDS

• Ionic bonds occur when atoms are attracted to each other by opposite electrical charges.• Strong (not as strong as covalent bonds)• Not directional• Form crystals

COVALENT BONDS

• Covalent bonds are formed when two atoms share electrons.• Atoms seek to fill outermost sphere of orbiting

electrons.• Strongest type of bond.• Very directional – bonds form between two specific

atoms rather than a generalized attraction of one atom for its neighbors.

HYDROGEN BONDS

• When covalent bonds form, one atom may attract the electrons more strongly than the other(s).

• Shared electrons spend more time around the “stronger” atom giving it a slightly negative charge, while the “weaker” atom(s) have a slightly positive charge.

HYDROGEN BONDS

• These charges are not as strong as those on an ion.

• The molecules end up with a positive end and a negative end and are said to be polar.

• A hydrogen bond occurs between the positive end of one polar molecule and the negative end of another polar molecule.

HYDROGEN BONDS

• Hydrogen bonds are weak.• As a result of weakness, they are highly

directional – polar molecules must be very close for the weak attraction to be effective.

• Act like Velcro, forming a tight bond as a result of many weak interactions.

HYDROGEN BONDS

• Hydrogen bonds are important in biology because they stabilize the shapes of biological molecules by causing certain parts to be attracted to other parts.

• Protein shape is dictated by hydrogen bonds.

HYDROGEN BONDS & WATER

• Water is essential for life.• H2O – An oxygen covalently bonded to two

hydrogens.• Water is polar – can form hydrogen bonds.• Weak, short-lived hydrogen bonds form

between water molecules.• Cumulative effect of transient bonds is

important.

HEAT STORAGE

• Temperature of a substance is a measure of how fast molecules are moving.

• Due to many hydrogen bonds, a large input of thermal energy is necessary to disrupt the organization of water and raise its temperature.

• Water heats slowly, holds temperature longer.

• Contributes to internal body temp regulation.

ICE FORMATION

• When temperature is low enough, few hydrogen bonds break.

• Water molecules form a crystal-like structure – ice.

• Ice is less dense than water – molecules are spaced out by hydrogen bonds.

HIGH HEAT OF VAPORIZATION

• When temperature is high enough, many hydrogen bonds break and liquid is changed to vapor.

• A great deal of heat energy is required to do this.

COHESION

• Polar water molecules are attracted to other polar molecules.

• Attraction between water molecules is called cohesion.

ADHESION

• When the other polar molecule is not water, the attraction is called adhesion.

• Capillary action is a result of adhesion.• Adhesion is why things get wet when

dipped in water.

HIGH POLARITY

• Water molecules tend to form the maximum number of hydrogen bonds possible.

• Polar molecules are called hydrophilic or water-loving molecules.

HIGH POLARITY

• When salt dissolves, molecules break off and are surrounded by water molecules.

• Polar molecules that dissolve in water are said to be soluble.

HIGH POLARITY

• Nonpolar molecules like oil do not form hydrogen bonds and are not water-soluble.

• When nonpolar substances are put in water, the water forms hydrogen bonds with each other – leaving out the nonpolar molecules.

• Nonpolar substances are said to be hydrophobic or water-fearing.

WATER IONIZES

• Covalent bonds within a water molecule sometimes break spontaneously.

H2O +OH–

hydroxide ion

H+

hydrogen ion

• This process of spontaneous ion formation is called ionization.• It is not common because of the strength of

covalent bonds.

PH

• A convenient way to express the hydrogen ion concentration of a solution.

pH= log [H+]_

• The pH scale is logarithmic• A difference of one unit represents a ten-fold

change in H+ concentration.

1

0

2

3

4

5

6

7

8

9

10

11

12

13

14 Sodium hydroxide

Examples ofSolutions

Hydrochloric acid

Stomach acid Lemon juice

Vinegar, cola, beer

Seawater

Baking soda

Great Salt LakeMilk of magnesia

Tomatoes

Black coffee Normal rain water

UrineSaliva

Pure waterBlood

Household ammonia

Household bleach

Oven cleaner

100

10-1

10-2

10-3

10-4

10-5

10-6

10-7

10-8

10-9

10-10

10-11

10-12

10-13

10-14

H+ IonConcentration

H+

H+

H+

H+H+

OH–

H+

H+ H+

H+H+

H+

H+

pH Value

OH–

OH–

OH–

OH–

OH–

OH–

OH–

OH–

OH–

OH–

PH

• Pure water has a pH of 7.• There are equal amounts of [H+] relative to

[OH–].• Acid—any substance that dissociates in

water and increases the [H+].• acidic solutions have pH values below 7.

• Base—any substance that combines with [H+] when dissolved in water.• basic solutions have pH values above 7.

PH

• The pH in most living cells and their environments is fairly close to 7.• Proteins involved in metabolism are

sensitive to any pH changes.• Acids and bases are routinely encountered by living organisms.• From metabolic activities (i.e., chemical

reactions).• From dietary intake and processing.

• Organisms use buffers to minimize pH disturbances.

WATER IONIZES

• Buffer—a chemical substance that takes up or releases hydrogen ions.• Buffers don’t remove the acid or the base

affecting pH but minimize their effect on it.• Most buffers are pairs of substances, one an

acid and one a base.

BUFFERS

• Living cells have a pH near 7.• A constant pH must be maintained for

metabolic activities to work properly (homeostasis).

• Buffers act as a reservoir for hydrogen ions, donating them or taking them as needed.

BUFFERS

• Key buffering pair in human blood: carbonic acid and bicarbonate (a base).