Transcript The Nature of Molecules chapt02_lecture

The Nature of Molecules
Chapter 2
Recall from Chapter 1:
Levels of Organization
Cellular Organization
cells
organelles
molecules
atoms
The cell is the
basic unit of life.
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Fig. 1.1-1
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Atomic Structure
• All matter is composed of atoms
• Understanding the structure of atoms is
critical to understanding the nature of
biological molecules:
• Bonding
• Reactivity
• Oxidation/ Reduction
• pH
• Energy
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Atomic Structure
• Atoms are composed of:
1. Protons – positively charged particles
2. Neutrons – neutral particles
3. Electrons – negatively charged particles
• Protons and neutrons are located in the
nucleus of the atom
• Electrons are found in orbitals
surrounding the nucleus
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Atomic Structure
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Atomic Structure
• Every different atom has a characteristic
number of protons in the nucleus.
• Atomic Number = number of protons
• Atoms with the same atomic number have
the same chemical properties and belong to
the same element.
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Atomic Structure
• Each proton and neutron has a mass of
approximately 1 dalton.
• The sum of protons and neutrons is the
atom’s atomic mass.
• Isotopes – atoms of the same element
that have different atomic mass numbers
due to different numbers of neutrons.
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Atomic Structure:
Carbon Isotopes
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Atomic Structure
• Neutral atoms have the same number of
protons and electrons
• Ions are charged atoms
•
Cations – have more protons than electrons and
are positively charged (+)
•
Anions – have more electrons than protons and
are negatively charged (-)
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Atomic Structure
Atomic Energy Levels
• Electrons have discrete energy levels
around the nucleus - Energy of Position
• Energy Levels are labelled K-M
• The distance an electron is from the
nucleus is relative to the amount of potential
energy of the electron
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Fig. 2.5
• Electrons release energy
as they fall from a high
energy level to a lower
energy level
• Electrons climb from a
lower energy level to a
higher energy level as
they absorb energy
Atomic Structure
Electron Orbitals
• At each Energy Level, electrons are
located in orbitals described as s or p
orbitals
• Orbitals are areas where the electron is
most likely to be found
• Each orbital can contain only 2 electrons
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Atomic Structure
Electron Orbitals
• At the K Energy Level, there is one s orbital
(referred to as the 1s orbital)
• There can be two electrons in the K Energy
Level
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Fig. 2.4a
Atomic Structure
Electron Orbitals
• At the L Energy Level, there is one s
orbital (referred to as the 2s orbital) and 3p
orbitals (each referred to as 2p orbitals)
• The L Energy Level holds 8 electrons
Fig. 2.4b
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Atomic StructureElectron
Configuration for Neon
• Neon has 10 electrons
• The K and L Energy Levels are
completely filled
• The 1s orbital of K and the 2s and three
2p orbitals of L are completely filled
Fig. 2.4c
Fig. 2.4
Atomic Structure
• Valence electrons are the electrons in the
outermost energy level of an atom
• Any level (K-M) can be the outermost level, it
depends on how many electrons the atom has
• An element’s chemical properties depend on
interactions between valence electrons of
different atoms
• These chemical properties include:
• Reactivity - atom’s ability to make or break
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bonds
Fig. 2.2
Atomic Structure
• Electrons can be transferred from one
atom to another, while still retaining the
energy of their position in the atom
• Oxidation = loss of an electron
• Reduction = gain of an electron
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Atomic Structure
LEO the lion says GER
If it Looses Electrons, it’s Oxidized
If it Gains Electrons, it’s Reduced
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Atomic Structure
• Oxidation and Reduction reactions always
occur in pairs
• If a an atom or molecule is reduced,
another atom or molecule must have been
oxidized
• If a an atom or molecule is oxidizied,
another atom or molecule must have been
reduced
• For this reason Oxidation and Reduction
Reactions are known as Redox Reactions
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Atomic Structure
Redox Reactions
• Some Terminology:
• The entity that lost (donates) the electrons in
a redox rxn. is oxidized
• Therefore, it is known as the Reducing Agent
in the Redox Rxn.
• It caused the other entity to be reduced by
providing the electron
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Atomic Structure
Redox Reactions
• Conversely:
•The entity that gained (accepted) the electrons
in a redox rxn. is reduced
• Therefore, it is known as the Oxidizing Agent
in the Redox Rxn.
• It caused the other entity to be Oxidized by
accepting the electron
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Atomic Structure
Redox Reactions
• Differentiate Oxidation/ Reduction and
•
•
•
Oxidizing Agent/ Reducing Agent
Remember, LEO says GER describes
Oxidation/ Reduction
The Oxidizing Agent is Reduced, The
Reducing Agent is Oxidized
Redox Reactions are very important in
biological processes (and easy test
questions)
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Elements
• The Periodic Table arranges all elements
according to their atomic number and their
number of Valence Electrons
• Therefore, the Periodic Table identifies and
groups elements with similar chemical
properties
• The number of Valence Electrons of an
element can be determined by counting
across the columns of the Periodic Table
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Periodic Table of the Elements
H has 1 e-
O has 6
Valence e-s
C has 4
Valence e-s
Na has 1
Valence e-
Ca has 2
Valence e-s
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Elements
• There are 90 naturally occurring elements
• Only 12 elements are found in living
organisms in substantial amounts
• Twelve or so others are found in trace
amounts: I, for example
• Four elements make up 96.3% of human
body weight: carbon, hydrogen, oxygen,
nitrogen
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Elements
• Become Familiar with the Valence Electron
Configuration for the Elements:
Hydrogen (H): 1
Carbon (C): 4
Oxygen (O): 6
Nitrogen (N): 5
Chlorine (Cl): 7
Sodium (Na): 1
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Elements
• Octet rule: States that Atoms typically
attempt to achieve 8 valence electrons
(H and He are exceptions)
• Eight valence electrons will fill the outer energy
level (H needs 1e-, He needs 2 e-s)
• Atoms will steal, give up, or share valence
electrons in order to fulfill the Octet Rule
• Atoms with full energy levels are more
stable and less reactive than atoms with
unfilled energy levels
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Elements
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Chemical Bonds
• Atoms can fulfill the octet rule by loosing,
gaining or sharing electrons with other
atoms, creating chemical bonds
• Atoms are held together in molecules or
compounds by chemical bonds
•
Molecules are groups of atoms held
together in a stable association
• Compounds are molecules containing more
than one type of element
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Chemical Bonds
• We will study 3 types of chemical bonds
1. Ionic Bonds
2. Covalent Bonds
3. Hydrogen Bonds
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Chemical Bonds
1. Ionic bonds
• Ionic Bonds are formed by the attraction
of oppositely charged ions
• Cation Na+ and Anion Cl- attract
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Chemical Bonds
1. Ionic bonds
• Na easily gives up its single valence
electron to fulfill the octet rule
• Because Na lost an electron it now has a
positive (+) charge
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Chemical Bonds
1. Ionic bonds
• Cl, with 7 valence electrons, easily picks
up the extra electron to fulfill the octet rule
• Because Cl gained an electron it now has
a positive (-) charge
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Chemical Bonds
1. Ionic bonds
• Cation Na+ and Anion Cl- attract
• Ionic Bonds form crystals
• Ionic Bonds dissociate in water
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Chemical Bonds
2. Covalent Bonds
• Covalent bonds form when atoms share 2 or
more valence electrons
• Covalent bond strength and length depends
on the number of electron pairs shared by the
atoms
• Single Covalent Bond - one pair of e-’s is shared
• Double Covalent Bond - two pairs of e-’s are shared
• Triple Covalent Bond - three pairs of e-’s are shared
• Covalent Bonds are the strongest bonds and
do not dissociate in water
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Chemical Bonds
2. Covalent Bonds
• Covalent bond strength depends on the
number of electron pairs shared by the atoms
• Single Covalent Bond - one pair of e-’s is shared
C
C
• Double Covalent Bond - two pairs of e-’s are shared
C
C
• Triple Covalent Bond - three pairs of e-’s are shared
C
C
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Chemical Bonds
2. Covalent Bonds
• Covalent bond strength depends on the
number of electron pairs shared by the atoms
• Single Covalent Bond - one pair of e-’s is shared
• Double Covalent Bond - two pairs of e-’s are shared
• Triple Covalent Bond - three pairs of e-’s are shared
single
• Bond Strength: bond
• Bond Length:
single
bond
<
double
bond
double
> bond
<
triple
bond
>
triple
bond
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Chemical Bonds
2. Covalent Bonds
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Chemical Bonds
• Electronegativity is an atom’s affinity for
electrons
• Atoms of each of the elements differ in
electronegativity
• Atomic size, mass, nuclear charge, electron
configuration all attribute to the differences in
electronegativity
• Based on what you know about the Na atom,
do you think Na has a high or low
electronegativity?
• What about Cl?
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Periodic Table of Elements
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Chemical Bonds
• Electronegativity is an atom’s affinity for
electrons
• The Periodic Table of Elements discloses
an Electronegativity Trend
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Electronegativity(EN) Trend in the
Periodic Table of Elements
Increasing EN
Greatest
EN
Lowest
EN
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Table 2.2
Chemical Bonds
• A covalent bond is formed by a sharing of
electrons, but electrons are not always
shared equally
• Differences in electronegativity dictate how
electrons are distributed in covalent bonds
• In a covalent bond between a highly
electronegative atom and a weakly
electronegative atom, electrons will be drawn
more towards the highly electronegative side
of the bond
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Chemical Bonds
• The water molecule is an important
example of a covalent bond between a highly
electronegative atom and a weakly
electronegative atom
• As a result of this unequal sharing of
electrons, one side of the molecule (the
highly electronegative side) has a partial
negative charge and the other side has a
partial positive charge
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Fig. 2.11a
Chemical Bonds
• A covalent bond with an unequal sharing of
electrons is known as a polar covalent bond
-
Polar because it has two ‘poles’, a partial
positive pole and a partial negative pole
• A covalent bond with an equal sharing of
electrons is known as a nonpolar covalent
bond
• A Carbon to Carbon covalent bond is an
example of a nonpolar covalent bond
C
C
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Chemical Bonds
• A molecule with an unequal sharing of
electrons is known as a polar molecule
- Water is an important polar molecule
• A molecule with an equal sharing of
electrons is known as a nonpolar molecule
-
A Carbon to Carbon molecule (like an oil molecule)
is an example of a nonpolar molecule
• Polar molecules mix only with other polar
molecules, nonpolar molecules mix only with
other nonpolar molecules
- ‘Like dissolves like’
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Chemical Bonds
• Chemical reactions involve the formation
or breaking of chemical bonds
• The making and breaking of bonds
• Whether a chemical reaction occurs is
influenced by many factors:
• temperature
• concentration of reactants and products
• availability of a catalyst
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Chemical Bonds
• Chemical reactions are written with the
reactants first, followed by the products
6H2O + 6CO2
reactants
C6H12O6 + 6O2
products
• Chemical reactions are often reversible
C6H12O6 + 6O2
6CO2
6H2O +
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Water Chemistry
• All living organisms are dependent on
water
• The structure of water is the basis for its
unique properties
• The most important property of water is
the ability to form hydrogen bonds
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Water Chemistry
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Water Chemistry
3. Hydrogen Bonds
• Hydrogen bonds are weak attractions
between the partially negative oxygen of
one water molecule and the partially positive
hydrogen of a different water molecule.
• Hydrogen bonds can form between water
molecules or between water and another
charged molecule.
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Water Chemistry
3. Hydrogen Bonds
• Within a water molecule, the bonds between
oxygen and hydrogen are highly polar.
• Partial electrical charges develop:
– oxygen side is partially negative
– hydrogen side is partially positive
•The partially negative oxygen side of one
water molecule is weakly attracted to the
partially negative hydrogen side of another
water molecule to form a hydrogen bond
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Water Chemistry
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Water Chemistry
• The polarity of water causes it to be
cohesive and adhesive.
• Cohesion: water molecules stick to other
water molecules by hydrogen bonding
• Adhesion: water molecules stick to other
polar molecules by hydrogen bonding
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Water Chemistry
Surface Tension
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Water Chemistry
Capillary Action
Look for the meniscus
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Properties of Water
1. Water has a high specific heat
• A large amount of energy is required to change
the temperature of water.
2. Water has a high heat of vaporization
• The evaporation of water from a surface causes
cooling of that surface.
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Properties of Water
3. Solid water is less dense than liquid water
• Bodies of water freeze from the top down
4. Water is a good solvent
•Water dissolves polar molecules - “like dissolves
like”
•Water dissolves ionic molecules
•Water Soluble
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Properties of Water
Water dissolves ionic molecules
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Properties of Water
5. Hyrophobic Interactions
• Water organizes nonpolar molecules
•Polar molecules are hydrophilic: “water-loving”
•Nonpolar molecules are hydrophobic: “waterfearing”
•Water causes hydrophobic molecules to aggregate
or assume specific shapes
•“Like dissolves like”, oil and water don’t mix
•A Micelle
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Properties of Water
6. Water can form 3 ions:

OH-1 +
H+1
hydroxide ion hydrogen
H 2O
water
ion
H 2O
water
+
H+1 
H3O+1
hydronium ion
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Properties of Water
6. Water can form 3 ions:
1. Hydrogen Ion
H 2O
water

OH-1 +
H+1
hydroxide ion hydrogen
ion
• The hydrogen ion is a lone proton with a +1
charge
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Fig. 2.2
Properties of Water
6. Water can form 3 ions:
2. Hydroxide Ion
H 2O
water

OH-1 +
H+1
hydroxide ion hydrogen
ion
• The hydoxide ion retains the electron
yielding a -1 charge
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Fig. 2.2
Properties of Water
6. Water can form 3 ions:
3. Hydronium Ion
H 2O
water
+
H+1 
H3O+1
hydronium ion
• A water molecule can pick up a hydrogen
ion (a proton, H+1) to become a hydronium
ion with +1 charge
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Table 2.3
Acids and Bases
• Acid: a chemical that releases H+1 ions
• Base: a chemical that accepts H+1 ions
• The acidity/basicity of a solution is
measured in terms of the pH Scale
Is water an Acid or a Base?
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Acids and Bases
Is water an acid or a base?
H2O  OH-1
H2O
+
H+1 
+
H+1
H3O+1
• Water acts as an acid and a base
• Pure water exists as a balance of H2O,
OH-1, and H3O+1
• Water has a neutral pH
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Acids and Bases
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Acids and Bases
The pH Scale
• Hydrogen ion concentration [H+1] is the
basis of the pH scale
pH = -log [H+1]
or
pH =
1
log [H+1]
• Greater H+1 concentration = lower pH
(acidic)
• Lower H+1 concentration = higher pH
(basic)
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Acids and Bases
The pH Scale
• Hydrogen ion concentration [H+1] of pure
water is 10-7 moles [H+1] / L
pH =
1
log [H+]
pHwater =
1
log [10-7]
pHwater = 7
•The pH scale is based on the pH of pure
water
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Acids and Bases
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Acids and Bases
Acid: a chemical that releases H+1 ions
• Hydrochloric acid, HCl
HCl  H+1 + Cl-
• Sulfuric acid, H2SO4
H2SO4  2H+1 + SO4-2
• Carbonic acid, H2CO3 (Example from text)
H2CO3  H+1 + HCO3Bicarbonate ion
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Acids and Bases
Base: a chemical that accepts H+1 ions
• Sodium Hydroxide, NaOH
NaOH  Na+1 + OH-1
OH-1 + H+1
 H 2O
• Bicarbonate, HCO3HCO3- + H+  H2CO3
carbonic acid
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Acids and Bases
Acid-Base Pairs
• Acids and Bases come in pairs
H2CO3 
Carbonic acid
H+1 + HCO3Bicarbonate
HCO3- + H+  H2CO3
Bicarbonate
Carbonic acid
H2CO3
H+1 + HCO381
Acids and Bases
Buffer Systems
Buffer: a chemical that resists a change in pH
• Buffers accept/release H+1 as necessary to
keep pH constant
• Buffers accept hydrogen ions in acidic
solutions and release hydrogen ions in basic
solutions, minimize a change in hydrogen ion
concentration
• The key buffer in human blood is the carbonic
acid/ bicarbonate acid-base pair
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Acids and Bases
Most biological buffers consist of a pair of
molecules, one an acid and one a base.
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Acids and Bases
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Table 2.1
The Nature of Molecules
End Chapter 2