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Transcript
CHAPTER # 2(a)
CHEMISTRY
COMES ALIVE
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
Energy
• Capacity to do work or put matter into motion
• Types of energy:
• Kinetic—energy in action
• Potential—stored (inactive) energy
PLAY
Animation: Energy Concepts
Copyright © 2010 Pearson Education, Inc.
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)
Copyright © 2010 Pearson Education, Inc.
Energy Form Conversions
• Energy may be converted from one form to
another
• Conversion is inefficient because some
energy is “lost” as heat
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
Composition of Matter
• Atoms
• Unique building blocks for each element
• Atomic symbol: one- or two-letter chemical
shorthand for each element
Copyright © 2010 Pearson Education, Inc.
Major Elements of the Human Body
• Oxygen (O)
• Carbon (C)
• Hydrogen (H)
• Nitrogen (N)
Copyright © 2010 Pearson Education, Inc.
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)
Copyright © 2010 Pearson Education, Inc.
Trace Elements of the Human Body
• < 0.01% of body mass:
• Part of enzymes, e.g., chromium (Cr),
manganese (Mn), and zinc (Zn)
Copyright © 2010 Pearson Education, Inc.
Atomic Structure
• Determined by numbers of subatomic
particles
• Nucleus consists of neutrons and protons
Copyright © 2010 Pearson Education, Inc.
Atomic Structure
• Neutrons
• No charge
• Mass = 1 atomic mass unit (amu)
• Protons
• Positive charge
• Mass = 1 amu
Copyright © 2010 Pearson Education, Inc.
Atomic Structure
• Electrons
• Orbit nucleus
• Equal in number to protons in atom
• Negative charge
• 1/2000 the mass of a proton (0 amu)
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
Models of the Atom
• Planetary model—oversimplified, outdated
model
• Incorrectly depicts fixed circular electron paths
• Useful for illustrations (as in the text)
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
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)
Copyright © 2010 Pearson Education, Inc.
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
Identifying Elements
• Atomic number = number of protons in
nucleus
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
Identifying Elements
• Atomic weight = average of mass numbers of
all isotopes
Copyright © 2010 Pearson Education, Inc.
Proton
Neutron
Electron
Hydrogen (1H)
(1p+; 0n0; 1e–)
Copyright © 2010 Pearson Education, Inc.
Deuterium (2H)
(1p+; 1n0; 1e–)
Tritium (3H)
(1p+; 2n0; 1e–)
Figure 2.3
Radioisotopes
• Spontaneous decay (radioactivity)
• Similar chemistry to stable isotopes
• Can be detected with scanners
Copyright © 2010 Pearson Education, Inc.
Radioisotopes
• Valuable tools for biological research and
medicine
• Cause damage to living tissue:
• Useful against localized cancers
• Radon from uranium decay causes lung
cancer
Copyright © 2010 Pearson Education, Inc.
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
atoms bonded together (e.g., C6H12O6)
Copyright © 2010 Pearson Education, Inc.
Mixtures
• Most matter exists as mixtures
• Two or more components physically
intermixed
• Three types of mixtures
• Solutions
• Colloids
• Suspensions
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
Concentration of Solutions
• Expressed as
• Percent, or parts per 100 parts
• Milligrams per deciliter (mg/dl)
• Molarity, or moles per liter (M)
• 1 mole = the atomic weight of an element or
molecular weight (sum of atomic weights) of
a compound in grams
• 1 mole of any substance contains 6.02 
1023 molecules (Avogadro’s number)
Copyright © 2010 Pearson Education, Inc.
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, e.g., blood
• Large visible solutes tend to settle out
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
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)
Copyright © 2010 Pearson Education, Inc.
Chemically Inert Elements
• Stable and unreactive
• Outermost energy level fully occupied or
contains eight electrons
Copyright © 2010 Pearson Education, Inc.
(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
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
Copyright © 2010 Pearson Education, Inc.
(b)
Chemically reactive elements
Outermost energy level (valence shell) incomplete
1e
Hydrogen (H)
(1p+; 0n0; 1e–)
6e
2e
Oxygen (O)
(8p+; 8n0; 8e–)
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
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
Formation of an Ionic Bond
• Ionic compounds form crystals instead of
individual molecules
• NaCl (sodium chloride)
Copyright © 2010 Pearson Education, Inc.
CI–
Na+
(c) Large numbers of Na+ and Cl– ions
associate to form salt (NaCl) crystals.
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
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
Covalent Bonds
• Sharing of electrons may be equal or unequal
• Equal sharing produces electrically balanced
nonpolar molecules
• CO2
Copyright © 2010 Pearson Education, Inc.
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
Copyright © 2010 Pearson Education, Inc.
Figure 2.8b
Copyright © 2010 Pearson Education, Inc.
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
PLAY
Animation: Hydrogen Bonds
Copyright © 2010 Pearson Education, Inc.
+
–
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
(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.
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
Examples of Chemical Equations
H + H  H2 (hydrogen gas)
(reactants)
(product)
4H + C  CH4 (methane)
Copyright © 2010 Pearson Education, Inc.
Patterns of Chemical Reactions
• Synthesis (combination) reactions
• Decomposition reactions
• Exchange reactions
Copyright © 2010 Pearson Education, Inc.
Synthesis Reactions
• A + B  AB
• Always involve bond formation
• Anabolic
Copyright © 2010 Pearson Education, Inc.
(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
Decomposition Reactions
• AB  A + B
• Reverse synthesis reactions
• Involve breaking of bonds
• Catabolic
Copyright © 2010 Pearson Education, Inc.
(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
Copyright © 2010 Pearson Education, Inc.
Figure 2.11b
Exchange Reactions
• AB + C  AC + B
• Also called displacement reactions
• Bonds are both made and broken
Copyright © 2010 Pearson Education, Inc.
(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
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
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.
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
Copyright © 2010 Pearson Education, Inc.