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Chapter 2:

Atoms, Molecules & LifeWhat Are Atoms?

• An atom are the smallest unit of matter.

• Atoms are composed of

�Electrons = negatively charged particles.

�Neutrons = particles with no charge (neutral).

�Protons = positively charged particles.

• Protons and neutrons are found in the nucleus of an atom

• Electrons orbit the nucleus.

�An atom is neutral, the # of electrons = # protons

An example of an atom:

Remember that neutrons are neutral, so there can be more of them than protons.

Periodic Table of Elements:

Element = Substance that can’t be

broken down or converted to another, simpler substance by ordinary chemical means.

Chapter 2: Atoms, Molecules & Life

Atomic Number = Number of protons in

the nucleus

� Since # of protons always equals the # of electrons, why don’t we use the # of electrons as the atomic number?

Atomic Mass = Number of protons &

neutrons in nucleus

Periodic Table

Why is the mass number a decimal? Isotope = The same element with a differentnumber of neutrons

• C14 • P32 • U235

Chapter 2: Atoms, Molecules & Life

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Isotopes

• Isotopes: Same element,

different number of neutrons.

• Different number of neutrons

changes the atomic mass, but

NOT the atomic number.

Atomic number remains 1 for hydrogen and its

isotopes

Isotopes

• Some isotopes, but not all, are radioactive.

�Example: Carbon 14 (C14) is radioactive

�Example: Hydrogen 2 (not radioactive) and hydrogen 3

(radiocative)

• Isotopes are useful in research

�Nuclear experiments involved heavy water (H2)

�Radiolabelling

�used to be H3, but now other isotopes are used.

Radiocarbon Dating:

• Technique for determining the age of materials that contain carbon based on C14 levels

Half-life ~ 5730 years

C14O2C12O2

1 C14 for every

1,000,000,000,000 C12

Crucial elements in life

• Carbon

�All organic matter has carbon. �18.5% of the human body mass is carbon atoms.

• Hydrogen

�All macromolecules have hydrogen as a component.�9.5% of the human body mass is hydrogen atoms.

• Oxygen

�All macromolecules have oxygen as a component.�65% of the human body mass is oxygen atoms. Why so much oxygen?

• Nitrogen

�All proteins have nitrogen as a component�3.3% of the human body mass is nitrogen atoms (mainly in muscle and

other proteins)

Other important elements in life

• Calcium

�Component of bones.

• Phosphorus

�A component of all cells (phospholipids).

• Potassium

�An important electrolyte, also keeps cell alive via sodium potassium pump.

• Sulfur

�A component of some protein molecules.

• Sodium

�Another important electrolyte, sodium ion pumps.

Compounds vs. Molecules

• Compound: A substance made up of different

types of atoms.

�Example: Table salt, NaCl.

• Molecule: a particle composed of one or more

atoms held together by chemical bonds.

�Example: Table salt, NaCl

�Also the smallest unit of a compound.

• Not all molecules are compounds.

�H2, O2, and other diatomic gases are not molecules.

Why?

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How are molecules formed?

• The number of electrons in the outermost electron

shell determine whether an atom is reactive or inert.

Carbon: reactive.

4 electrons in outer shell, needs 4 more to fill shell

Neon: inert

8 electrons in outer shell, Does not need electrons to fill shell

How are molecules formed?

• How many electrons are

needed to fill an electron

shell?

�Depends on which shell.

• First shell only needs two

�hydrogen, helium.

• Second shell needs eight.

• Third shell needs eight.

Inert Atoms = Atoms with their outermost shell either completely full or empty

Reactive Atoms = Atoms with their outermost shell only partially filled

Helium

Neon

Argon

Hydrogen

Carbon

Oxygen

How are molecules formed?

• Reactive atoms want to lose or gain electrons to stabilize their outer (valence) shell.

(Figure 2.3)

Chemical bonds

• Types of chemical bonds

�Covalent: the strongest of the three main types.

�Ionic: weaker than covalent bonds.

�Hydrogen bonds: the weakest chemical bond of the

three main types.

• Chemical bonds are crucial for chemical reactions

�Chemical Reaction = making or breaking chemical

bonds.

�Chemical reactions are essential for all life.

Covalent bonds: the strongest chemical bond

• Atoms share electrons in covalent bonds.

Sharing of

electrons

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Covalent bonds: the strongest

• Most biological molecules utilize covalent bonds.

�Proteins

�Carbohydrates

�Lipids

�DNA

• Carbon atoms are always linked by covalent bonds.

Examples of Covalent Bonds:

H - H

Share one pair of electrons (H2)

• Single covalent bond

Share two pairs of electrons (O2)

• Double covalent bond

O = O

Share three pairs of electrons (N2)

• Triple covalent bondN N

Bonding Patterns (Table 2.3):

Carbon (C)

Oxygen (O)

Non-polar Covalent Bonds:

Equal sharing of electrons

Polar Covalent Bonds:

Unequal sharing of electrons

Molecule is electrically neutral, but

poles are charged due to differences

in nuclear attraction for electrons

(electronegativity)

Ionic Bonds

• Attractive force between atoms that have lost or gained electrons

�Creates ions (negatively or positively charged atoms)

Transfer of

electrons

Example of ionic bonds

• Sodium chloride (table salt)

�Sodium wants to lose one electron to stabilize its outer shell

�Chloride needs to gain one electron to stabilize its outer shell.

Sodium Chloride (NaCl - Figure 2.3)

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Ionic bonds typically occur between atoms that are located on opposite sides of the periodic table.

Hydrogen Bonds

• Attractive force between water molecules due to polar covalent bonds.

�Electrons are far more attracted to oxygen than to hydrogen atoms.

�hydrogen atoms have a slight positive charge.

�Oxygen atoms have a slight negative charge.

• Makes water a very special molecule.

(Figure 2.5)

Water = Good Stuff!

Why is Water so Important to Life?

Life most likely arose in water

Living organisms 60 - 90% water

Humans and water

• Human body is 65% water.

• The average human can

survive several months

without food.

�But you can only survive

3 – 4 days without water.

All life depends on water

• Search for life on other planets often includes

the search for H2O

Why is Water so Important to Life?

• Phoenix spacecraft will land in one

of Martian ice caps in hopes of finding water and microbial life.

� Landed May, 2008, found ice on

July 31st, 2008.

1) Water is an excellent solvent:

Importance of Water:

• liquid capable of dissolving other substances in itself

Dissolving Ionic Bonds:(Salt)

(Figure 2.6)

Polar nature of water

Solution = Fluid containing dissolved substances

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• Molecules electricallyattracted to water

• Ions, polar molecules

Hydrophilic Molecules:

Water-loving

Hydrophobic Molecules:

Water-fearing

• Molecules electricallyneutral (fats / oils)

• Molecules tend to clumptogether in water(Figure 2.7)

Dissolving Polar Covalent Bonds:(e.g. glucose)

+-

Hydrophilic vs hydrophobic

Hydrophilic Hydrophobic

2) Water molecules tend to stick together (cohesion):

Importance of Water:

• Surface Tension: Tendency for a water surface to resist breaking

Walk on

water

Flow against

gravity

• Adhesion: Tendency for water to stick to walls of surfaces

3) Water Can Form H+ and OH- Ions (ionization):

Importance of Water:

Pure water contains equal amounts of H+ and OH-

HCl Acidic

H+ = OH-

Water

H+ = OH-

Water

NaOHBasic

OH-H+ <H+ > OH-

H2O H+ OH-+

Acid

A substance that increases the [H+]

in a solution

Base

A substance that decreases the

[H+] in a solution

Acids & bases disrupt the equilibrium. The pH of a solution describes its degree of acidity:

(Figure 2.9)

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Importance of pH to living systems

Changes in pH disrupt life chemistry

Example: Blood

usual pH = 7.4

pH 7.0 , 7.8 lethal

So how does blood maintain a healthy pH?

Buffers

Buffers - solutes that act to resist changes to

the pH of a solution when H+ or OH- is

added.

Biological fluids use buffers to help maintain

correct pH.

Buffers maintain a solution at relatively constant pH:

• Buffers either accept or release a H+ in response to changes in pH

• Stable pH essential for normal function

Example: Bicarbonate ion (HCO3-)

Too acidic?

H+

Hydrogen ion

+HCO3-

Bicarbonate ion

H2CO3Carbonic Acid

Too basic?

OH-

Hydroxide ion

+H2CO3Carbonic Acid

H20HCO3- +

Bicarbonate ion Water

4) Water Moderates Temperature Changes:

Importance of Water:

• Slow molecules = Cool temperatures

• Fast molecules = Warm temperatures

• Temperature = Speed of molecules

Background:

A) Water Heats Slowly

• Energy first initiates breaking of hydrogen bonds…

• Specific Heat = energy needed to heat 1 gram of

a substance 1°C

• Specific HeatWater = 1 cal

• Specific HeatAlcohol = 0.6 cal

• Specific HeatGranite = 0.02 cal

4) Water Moderates Temperature Changes:

Importance of Water:

A) Water heats slowly

B) Water is an effective coolant

• Heat of Vaporization: Heat needed to convert liquidwater to water vapor

• 529 calories/gram (very high!)

By evaporating 1 g of water,

539 grams of human body cools 1°C

C) Water freezes slowly

• Moderates the effects of low temperatures

D) Water forms ice (less dense than fluid water):

• Acts as an insulator for life below

Why does ice weigh less than liquid water?

• Water is the most dense (0.9998 g/cm3) at 40 C.

• Below 4 degrees, the water molecules form more hydrogen bonds, forming lattice like crystals.

• This characteristic is very important for aquatic organisms.

�Why??

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Because ice floats!

Any questions?

Next Class

• We will be starting Chapter 3: Biological

molecules

• Remember there are games to use as study tools

�Especially useful for learning terminology.