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Chapter 2:
Atoms, Molecules & Life
What 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.
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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
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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
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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 ofelectrons
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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)
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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
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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
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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
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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 liquid
water 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.
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