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Part A
Chemistry Comes Alive
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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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Capacity to do work or put matter into
motion
Types of energy:
◦ Kinetic—energy in action
◦ Potential—stored (inactive) energy
PLAY
Animation: Energy Concepts
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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 can neither be created nor destroyed
(1st law of thermodynamics)
Energy may be converted from one form to
another
Conversion is inefficient because some
energy is “lost” as heat
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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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Atoms
◦ Unique building blocks for each element
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Atomic symbol: one- or two-letter chemical
shorthand for each element
Eg: Copper
Iron
Mercury
Gold
Cu
Fe
Hg
Au
Sulfur
Potassium
Phosphorus
Iodine
S
K
P
I
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Oxygen (O)
Carbon (C)
Hydrogen (H)
Nitrogen (N)
About 96% of body mass
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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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< 0.01% of body mass:
◦ Part of enzymes, e.g., chromium (Cr), manganese
(Mn), and zinc (Zn)
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Determined by numbers of subatomic
particles
Nucleus consists of neutrons and protons
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Neutrons
 No charge
 Mass = 1 atomic mass unit (amu)
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Protons
 Positive charge
 Mass = 1 amu
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Electrons
◦
◦
◦
◦
Orbit nucleus
Equal in number to protons in atom
Negative charge
1/2000 the mass of a proton (0 amu) Not
considered in the calculation of atomic weight
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Planetary model
◦ Depicts fixed circular electron paths
◦ Useful for illustrations (as in the text)
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
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Atoms of different elements contain different
numbers of subatomic particles
◦ Compare hydrogen, helium and lithium (next slide)
Proton
Neutron
Electron
Hydrogen (H)
(1p+; 0n0; 1e–)
Copyright © 2010 Pearson Education, Inc.
Helium (He)
(2p+; 2n0; 2e–)
Lithium (Li)
(3p+; 4n0; 3e–)
Figure 2.2
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Atomic number = number of protons in
nucleus
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Atomic weight = mass of the protons and
neutrons
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
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Spontaneous decay (radioactivity)
Similar chemistry to stable isotopes
Can be detected with scanners
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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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Most atoms combine chemically with other
atoms to form molecules and compounds
◦ Molecule—two or more atoms of same element
bonded together (e.g., H + H = H2 )
◦ Compound—two or more atoms of different
elements bonded together (e.g., C6H12O6)
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Most matter exists as mixtures
◦ Two or more components physically intermixed
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Three types of mixtures
◦ Solutions
◦ Colloids
◦ Suspensions
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Solutions are homogeneous mixtures
Usually transparent, e.g., atmospheric air or
seawater
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◦ Solvent
 Present in greatest amount, usually a liquid
◦ Solute(s)
 Present in smaller amounts
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Colloids (emulsions)
◦ Heterogeneous translucent mixtures, e.g., cytosol
◦ Large solute particles that do not settle out
◦ Undergo sol-gel transformations
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Suspensions:
◦ Heterogeneous mixtures, e.g., blood
◦ Large visible solutes tend to settle out
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
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Mixtures
◦ No chemical bonding between components
◦ Can be separated physically, such as by straining or
filtering
◦ Heterogeneous or homogeneous
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Compounds
◦ Can be separated only by breaking bonds
◦ All are homogeneous
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Electrons occupy up to seven electron shells
(energy levels) around nucleus
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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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Stable and unreactive
Outermost energy level fully occupied or
contains eight electrons
(a)
Chemically inert elements
Outermost energy level (valence shell) complete
8e
2e
Helium (He)
(2p+; 2n0; 2e–)
Copyright © 2010 Pearson Education, Inc.
2e
Neon (Ne)
(10p+; 10n0; 10e–)
Figure 2.5a
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Outermost energy level not fully occupied by
electrons
Tend to gain, lose, or share electrons (form
bonds) with other atoms to achieve stability
(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
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Ionic
Covalent
Hydrogen
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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
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Attraction of opposite charges results in an
ionic bond
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.
Copyright © 2010 Pearson Education, Inc.
(b) After electron transfer, the oppositely
charged ions formed attract each other.
Figure 2.6a-b
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Ionic compounds form crystals instead of
individual molecules
◦ NaCl (sodium chloride)
CI–
Na+
(c) Large numbers of Na+ and Cl– ions
associate to form salt (NaCl) crystals.
Copyright © 2010 Pearson Education, Inc.
Figure 2.6c
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Formed by sharing of two or more valence
shell electrons
Allows each atom to fill its valence shell at
least part of the time
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.
Copyright © 2010 Pearson Education, Inc.
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.
Copyright © 2010 Pearson Education, Inc.
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.
Copyright © 2010 Pearson Education, Inc.
Structural
formula
shows
triple
bond.
Figure 2.7c
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Sharing of electrons may be equal or unequal
◦ Equal sharing produces electrically balanced
nonpolar molecules
 CO2
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Figure 2.8a
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Unequal sharing by atoms with different
electron-attracting abilities produces polar
molecules
◦ H 2O
 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
Copyright © 2010 Pearson Education, Inc.
Figure 2.8b
Copyright © 2010 Pearson Education, Inc.
Figure 2.9
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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
PLAY
Animation: Hydrogen Bonds
+
–
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.
Copyright © 2010 Pearson Education, Inc.
Figure 2.10a
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Occur when chemical bonds are formed,
rearranged, or broken
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Represented as chemical equations
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Chemical equations contain:
◦ Molecular formula for each reactant and product
◦ Relative amounts of reactants and products, which
should balance
H + H  H2 (hydrogen gas)
(reactants)
(product)
4H + C  CH4 (methane)
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Synthesis (combination) reactions
Decomposition reactions
Exchange reactions
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A + B  AB
◦ Always involve bond formation
◦ Anabolic
(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
Copyright © 2010 Pearson Education, Inc.
Figure 2.11a
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AB  A + B
◦ Reverse synthesis reactions
◦ Involve breaking of bonds
◦ Catabolic
(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
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AB + C = AC + B
AB + CD = AD + CB
◦ Also called displacement reactions
◦ Bonds are both made and broken
(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
Copyright © 2010 Pearson Education, Inc.
Adenosine diphosphate (ADP)
Figure 2.11c
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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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All chemical reactions are either exergonic or
endergonic
◦ Exergonic reactions—release energy
 Catabolic reactions
◦ Endergonic reactions—form energy bonds: products
contain more potential energy than did reactants
 Anabolic reactions
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All chemical reactions are theoretically
reversible
◦ A + B  AB
◦ AB  A + B
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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 reaction is influenced by:
◦  temperature   rate
◦  particle size   rate
◦  concentration of reactant   rate
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Catalysts:  rate without being chemically
changed: they reduce the activation energy
needed for the reaction
◦ Enzymes are biological catalysts