Transcript Chapter 2a

PowerPoint® Lecture Slides
prepared by
Janice Meeking,
Mount Royal College
CHAPTER
2
Chemistry
Comes Alive:
Part A
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Matter
•
Anything that has mass and occupies
space
•
States of matter:
1. Solid—definite shape and volume
2. Liquid—definite volume, changeable shape
3. Gas—changeable shape and volume
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Mass and Weight
• the mass of an object is a fundamental property
of the object
• a numerical measure of its inertia
• measure of the amount of matter in the object.
• definitions of mass often seem circular because it
is such a fundamental quantity that it is hard to
define in terms of something else
• the usual symbol for mass is m and its SI unit is
the kilogram
• the weight of an object is the force of gravity on
the object (w = mg)
3
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Energy
• Capacity to do work or put matter into motion
• Types of energy:
• Kinetic — energy in action
• Potential — stored (inactive) energy
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Energy Concepts
• What is energy?
•
The capacity to perform work
• What is the difference between potential and kinetic energy?
•
Stored vs. motion
• Energy is neither created nor destroyed but…
•
Converted from one form to another
•
This property is called the conservation of energy
• What is the usual way in which energy is “lost?”
•
Through heat
• What type of energy is heat?
•
Kinetic due to random motion of atoms
•
Heat is generated by friction (in this example between atoms and air)
• Heat is highly __________ energy and highest amount of _________.
•
Disordered, entropy
• Chemical energy is a form of ____________ energy.
•
Potential
• What is the primary form of chemical energy in living organisms?
•
ATP
• What is cellular respiration? What are the byproducts?
•
Conversion of glucose into ATP through reduction of oxygen forming water and carbon
dioxide
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Forms of Energy
• Chemical energy — stored in bonds of
chemical substances
• Electrical energy — results from movement of
charged particles
• Mechanical energy — directly involved in
moving matter
• Radiant or electromagnetic energy — exhibits
wavelike properties (i.e., visible light,
ultraviolet light, and X-rays)
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Energy Form Conversions
• Energy may be converted from one form to
another
• Conversion is inefficient because some
energy is “lost” as heat
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Composition of Matter
• Elements
• Cannot be broken down by ordinary chemical means
• Each has unique properties:
• Physical properties
• Are detectable with our senses or are
measurable
• Chemical properties
• How atoms interact (bond) with one another
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Composition of Matter
• Atoms
• Unique building blocks for each element
• Atomic symbol: one- or two-letter chemical
shorthand for each element
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Major Elements of the Human Body
• Oxygen (O)
• Carbon (C)
• Hydrogen (H)
• Nitrogen (N)
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About 96% of body mass
Lesser Elements of the Human Body
• About 3.9% of body mass:
• Calcium (Ca), phosphorus (P), potassium (K),
sulfur (S), sodium (Na), chlorine (Cl),
magnesium (Mg), iodine (I), and iron (Fe)
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Trace Elements of the Human Body
• < 0.01% of body mass:
• Part of enzymes, e.g., chromium (Cr),
manganese (Mn), and zinc (Zn)
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Atomic Structure
• Determined by numbers of subatomic
particles
• Nucleus consists of neutrons and protons
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Preferred web sites for atoms and structure
(note: the last 2 can only be accessed with
registration through your book)
• http://www.youtube.com/watch?v=zEX2aGpID
BY&feature=related
• http://www.youtube.com/watch?v=TBrJt5LHgQ&feature=related
• http://www.myaandp.com/booksAvail.html
• http://wps.aw.com/bc_marieb_hap_8/102/261
43/6692813.cw/index.html
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Atomic Structure
• Neutrons
• No charge
• Mass = 1 atomic mass unit (amu)
• Protons
• Positive charge
• Mass = 1 amu
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Atomic Structure
• Electrons
• Orbit nucleus
• Equal in number to protons in atom
• Negative charge
• 1/2000 the mass of a proton (0 amu)
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Models of the Atom
• Orbital model: current model used by
chemists
• Depicts probable regions of greatest electron
density (an electron cloud)
• Useful for predicting chemical behavior of
atoms
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Models of the Atom
• Planetary model — oversimplified, outdated
model
• Incorrectly depicts fixed circular electron paths
• Useful for illustrations (as in the text)
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Nucleus
Nucleus
Helium atom
Helium atom
2 protons (p+)
2 neutrons (n0)
2 electrons (e–)
2 protons (p+)
2 neutrons (n0)
2 electrons (e–)
(a) Planetary model
Proton
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Neutron
(b) Orbital model
Electron
Electron
cloud
Figure 2.1
Identifying Elements
• Atoms of different elements contain different
numbers of subatomic particles
• Compare hydrogen, helium and lithium (next
slide)
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Proton
Neutron
Electron
Hydrogen (H)
(1p+; 0n0; 1e–)
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Helium (He)
(2p+; 2n0; 2e–)
Lithium (Li)
(3p+; 4n0; 3e–)
Figure 2.2
Identifying Elements
• Atomic number = number of protons in
nucleus
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Identifying Elements
• Mass number = mass of the protons and
neutrons
• Mass numbers of atoms of an element are not
all identical
• Isotopes are structural variations of elements
that differ in the number of neutrons they
contain
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Identifying Elements
• Atomic weight = average of mass numbers of
all isotopes
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Biologically important atoms
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Radioisotopes
• Spontaneous decay (radioactivity)
• Similar chemistry to stable isotopes
• Can be detected with scanners
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Proton
Neutron
Electron
Hydrogen (1H)
(1p+; 0n0; 1e–)
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Deuterium (2H)
(1p+; 1n0; 1e–)
Tritium (3H)
(1p+; 2n0; 1e–)
Figure 2.3
Radioisotopes
• Valuable tools for biological research and
medicine
• Cause damage to living tissue:
• Useful against localized cancers
• Radon from uranium decay causes lung
cancer
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Molecules and Compounds
• Most atoms combine chemically with other
atoms to form molecules and compounds
• Molecule — two or more atoms bonded
together (e.g., H2 or C6H12O6)
• Compound — two or more different kinds of
elements bonded together (e.g., C6H12O6)
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Mixtures
• Most matter exists as mixtures
• Two or more components physically
intermixed
• Three types of mixtures
• Solutions
• Colloids
• Suspensions
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Solutions
• Homogeneous mixtures
• Usually transparent, e.g., atmospheric air or
seawater
• Solvent
• Present in greatest amount, usually a liquid
• Solute(s)
• Present in smaller amounts
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Colloids and Suspensions
• Colloids (emulsions)
• Heterogeneous translucent mixtures, e.g.,
cytosol
• Large solute particles that do not settle out
• Undergo sol-gel transformations
• Suspensions:
• Heterogeneous mixtures (blood)
• Large visible solutes tend to settle out
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Solution
Colloid
Suspension
Solute particles are very
tiny, do not settle out or
scatter light.
Solute particles are larger
than in a solution and scatter
light; do not settle out.
Solute particles are very
large, settle out, and may
scatter light.
Solute
particles
Solute
particles
Solute
particles
Example
Example
Example
Mineral water
Gelatin
Blood
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Figure 2.4
Mixtures vs. Compounds
• Mixtures
• No chemical bonding between components
• Can be separated physically, such as by
straining or filtering
• Heterogeneous or homogeneous
• Compounds
• Can be separated only by breaking bonds
• All are homogeneous
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Chemical Bonds
• Electrons occupy up to seven electron shells
(energy levels) around nucleus
• Octet rule: Except for the first shell which is
full with two electrons, atoms interact in a
manner to have eight electrons in their
outermost energy level (valence shell)
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Chemically Inert Elements
• Stable and unreactive
• Outermost energy level fully occupied or
contains eight electrons
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(a)
Chemically inert elements
Outermost energy level (valence shell) complete
8e
2e
Helium (He)
(2p+; 2n0; 2e–)
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2e
Neon (Ne)
(10p+; 10n0; 10e–)
Figure 2.5a
Chemically Reactive Elements
• Outermost energy level not fully occupied by
electrons
• Tend to gain, lose, or share electrons (form
bonds) with other atoms to achieve stability
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(b)
Chemically reactive elements
Outermost energy level (valence shell) incomplete
1e
Hydrogen (H)
(1p+; 0n0; 1e–)
6e
2e
Oxygen (O)
(8p+; 8n0; 8e–)
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4e
2e
Carbon (C)
(6p+; 6n0; 6e–)
1e
8e
2e
Sodium (Na)
(11p+; 12n0; 11e–)
Figure 2.5b
Types of Chemical Bonds
• Ionic
• Covalent
• Hydrogen
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Ionic Bonds
• Ions are formed by transfer of valence shell
electrons between atoms
• Anions (– charge) have gained one or more
electrons
• Cations (+ charge) have lost one or more
electrons
• Attraction of opposite charges results in an
ionic bond
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Sodium atom (Na)
(11p+; 12n0; 11e–)
Chlorine atom (Cl)
(17p+; 18n0; 17e–)
+
–
Sodium ion (Na+)
Chloride ion (Cl–)
Sodium chloride (NaCl)
(a) Sodium gains stability by losing one electron, and
chlorine becomes stable by gaining one electron.
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(b) After electron transfer, the oppositely
charged ions formed attract each other.
Figure 2.6a-b
Formation of an Ionic Bond
• Ionic compounds form crystals instead of
individual molecules
• NaCl (sodium chloride)
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CI–
Na+
(c) Large numbers of Na+ and Cl– ions
associate to form salt (NaCl) crystals.
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Figure 2.6c
Covalent Bonds
• Formed by sharing of two or more valence
shell electrons
• Allows each atom to fill its valence shell at
least part of the time
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Reacting atoms
Resulting molecules
+
Molecule of
Hydrogen
Carbon
methane gas (CH4)
atoms
atom
(a) Formation of four single covalent bonds:
carbon shares four electron pairs with four
hydrogen atoms.
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or
Structural
formula
shows
single
bonds.
Figure 2.7a
Reacting atoms
Resulting molecules
+
Oxygen
atom
or
Oxygen
atom
Molecule of
oxygen gas (O2)
(b) Formation of a double covalent bond: Two
oxygen atoms share two electron pairs.
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Structural
formula
shows
double
bond.
Figure 2.7b
Reacting atoms
Resulting molecules
+
Nitrogen
atom
or
Nitrogen
atom
Molecule of
nitrogen gas (N2)
(c) Formation of a triple covalent bond: Two
nitrogen atoms share three electron pairs.
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Structural
formula
shows
triple
bond.
Figure 2.7c
Covalent Bonds
• Sharing of electrons may be equal or unequal
• Equal sharing produces electrically balanced
nonpolar molecules
• CO2
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Figure 2.8a
Covalent Bonds
• Unequal sharing by atoms with different
electron-attracting abilities produces polar
molecules
• H2O
• Atoms with six or seven valence shell
electrons are electronegative, e.g., oxygen
• Atoms with one or two valence shell
electrons are electropositive, e.g., sodium
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Figure 2.8b
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Figure 2.9
Hydrogen Bonds
• Attractive force between electropositive
hydrogen of one molecule and an
electronegative atom of another molecule
• Common between dipoles such as water
• Also act as intramolecular bonds, holding a
large molecule in a three-dimensional shape
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Hydrogen bonds
•
The bonds of a water molecule represent ________ _______ type of bond. Also known as a ________.
•
•
Oxygen has a greater affinity for the electrons and is therefore more _____________. Whereas, hydrogen
has a lesser attraction for electrons is more _____________.
•
•
•
Negative, positive
The attraction between the negative oxygen end of one water compound to the positive hydrogen end of
another water represents a ___________ bond.
•
•
Electronegative, electropositive
The oxygen end of the molecule is therefore slightly more _________ and the hydrogen ends are slightly
more _________.
•
•
Polar covalent, dipole
Hydrogen
Hydrogen bonds are strong bonds. (T/F)
•
False
•
They are easily broken
Hydrogen bonds may inter- or intramolecular. (T/F)
•
True
• The unique properties of water are attributable to hydrogen bonds. Some of the properties
include….
•
Cohesion, high boiling point, why ice floats, high heat of vaporization, high heat capacity
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+
–
Hydrogen bond
(indicated by
dotted line)
+
+
–
–
–
+
+
+
–
(a) The slightly positive ends (+) of the water
molecules become aligned with the slightly
negative ends (–) of other water molecules.
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Figure 2.10a
(b) A water strider can walk on a pond because of the high
surface tension of water, a result of the combined
strength of its hydrogen bonds.
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Figure 2.10b
Chemical Reactions
• Occur when chemical bonds are formed,
rearranged, or broken
• Represented as chemical equations
• Chemical equations contain:
• Molecular formula for each reactant and
product
• Relative amounts of reactants and products,
which should balance
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Examples of Chemical Equations
H + H  H2 (hydrogen gas)
(reactants)
(product)
4H + C  CH4 (methane)
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Patterns of Chemical Reactions
• Synthesis (combination) reactions
• Decomposition reactions
• Exchange reactions
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Synthesis Reactions
• A + B  AB
• Always involve bond formation
• Anabolic
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(a) Synthesis reactions
Smaller particles are bonded
together to form larger,
more complex molecules.
Example
Amino acids are joined together to
form a protein molecule.
Amino acid
molecules
Protein
molecule
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Figure 2.11a
Dehydration Synthesis and Hydrolysis
• http://www.youtube.com/watch?v=b7TdWLNh
MtM&feature=player_detailpage
• http://www.phschool.com/science/biology_pla
ce/biocoach/bioprop/monomers.html
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Dehydration Synthesis and Hydrolysis
• What is dehydration synthesis?
• Removal of a water molecule to form a new covalent
bond
• What is hydrolysis?
• The addition of a water molecule to break a covalent
bond
• What is anabolism?
• Forming new bonds to build something bigger.
Requires energy (endergonic)
• What is catabolism?
• Breaking bonds to make something smaller. Large
molecules down to subunits.
• Releases energy (exergonic).
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Decomposition Reactions
• AB  A + B
• Reverse synthesis reactions
• Involve breaking of bonds
• Catabolic
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(b) Decomposition reactions
Bonds are broken in larger
molecules, resulting in smaller,
less complex molecules.
Example
Glycogen is broken down to release
glucose units.
Glycogen
Glucose
molecules
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Figure 2.11b
Exchange Reactions
• AB + C  AC + B
• Also called displacement reactions
• Bonds are both made and broken
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(c) Exchange reactions
Bonds are both made and broken
(also called displacement reactions).
Example
ATP transfers its terminal phosphate
group to glucose to form glucose-phosphate.
+
Glucose
Adenosine triphosphate (ATP)
+
Glucose
phosphate
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Adenosine diphosphate (ADP)
Figure 2.11c
Oxidation-Reduction (Redox) Reactions
• Decomposition reactions: Reactions in which
fuel is broken down for energy
• Also called exchange reactions because
electrons are exchanged or shared differently
• Electron donors lose electrons and are
oxidized
• Electron acceptors receive electrons and
become reduced
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Chemical Reactions
• All chemical reactions are either exergonic or
endergonic
• Exergonic reactions — release energy
• Catabolic reactions
• Endergonic reactions — products contain
more potential energy than did reactants
• Anabolic reactions
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Chemical Reactions
• All chemical reactions are theoretically reversible
• A + B  AB
• AB  A + B
• Chemical equilibrium occurs if neither a forward nor
reverse reaction is dominant
• Many biological reactions are essentially irreversible
due to
• Energy requirements
• Removal of products
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Rate of Chemical Reactions
• Rate of reaction is influenced by:
•  temperature   rate
•  particle size   rate
•  concentration of reactant   rate
• Catalysts:  rate without being chemically
changed
• Enzymes are biological catalysts
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