Transcript 2 - FacultyWeb

Chapter 2
Chemistry Comes
Alive
Part A
Shilla Chakrabarty, Ph.D.
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Matter And Energy
Matter is 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
Energy is the capacity to do work or put matter into motion
Types of energy:
• Kinetic - energy in action
• Potential - stored (inactive) energy
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Forms of Energy and Energy Conversions
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)
Energy 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
• All matter is composed of elements
• Elements cannot be broken down by ordinary chemical means
• Each element has unique physical and chemical properties:
 Physical properties are detectable with our senses, or are
measurable
 Chemical properties determine how atoms interact (bond)
with one another
• Of the 92 naturally occurring elements, only 25 are essential for
living organisms
• Atoms are the unique building blocks for each element
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Major Elements
Lesser Elements
Trace Elements
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Properties Of An Element Depend On The Structure Of
Its Atoms
 Each element consists of unique atoms.
 An atom is
 the smallest unit of matter that still retains the properties of an
element.
 composed of even smaller parts, called subatomic particles.
 Neutrons and protons: packed together to form a dense core,
the Atomic Nucleus, at the center of an atom.
 Neutrons: No charge; mass = 1 atomic mass unit (amu)
 Protons: Positive charge; mass = 1 amu
 Electrons:
 Form a cloud around the nucleus; equal in number to protons
in an atom
 Negative charge, negligible mass (1/2000th that of protons and
neutrons)
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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
 Planetary model — oversimplified, outdated model
• Incorrectly depicts fixed circular electron paths
• Useful mostly for illustrations
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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
Atomic Number and Atomic Mass
• Atoms of the various elements differ in number of subatomic
particles
• Atomic number:
 Number of protons in the nucleus of an atom.
 Written as a subscript before the symbol for the element (example,
2He).
• Atoms generally have equal numbers of protons and
electrons - no net charge.
• Mass number:
 Is the sum of protons and neutrons in the nucleus of an atom
 Written as a superscript before an element’s symbol (example, 4He).
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Identifying Elements
Proton
Neutron
Electron
Atoms of different elements contain
different numbers of subatomic particles
Hydrogen (H)
(1p+; 0n0; 1e–)
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Helium (He)
(2p+; 2n0; 2e–)
Lithium (Li)
(3p+; 4n0; 3e–)
Figure 2.2
Isotopes Of Elements
All atoms of an element have the same number of protons
(atomic number) but may differ in number of neutrons
Proton
Neutron
Electron
Isotopes are two atoms of an element that differ in number
of neutrons.
In nature, an element occurs as a mixture of isotopes.
Hydrogen (1H)
(1p+; 0n0; 1e–)
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Deuterium (2H)
(1p+; 1n0; 1e–)
Tritium (3H)
(1p+; 2n0; 1e–)
Radioisotopes And Their Applications
• Radioactive isotopes
• Decay spontaneously, giving off particles and energy
Example: 14C is a radioactive isotope. In its decay, a neutron is
converted to a proton and an electron. This converts 14C to 14N,
changing the identity of that atom.
• Cause damage to living tissues
Example: Radon from Uranium decay causes Lung Cancer.
• Can be detected by scanners
• Some applications of radioactive isotopes in biological research:
 Dating fossils
 Tracing atoms through metabolic processes
 Diagnosing and treating medical conditions
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Molecules, Compounds and Mixtures
• 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)
• Most matter exists as mixtures
 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
 Concentration of solutions is 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)
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Colloids and Suspensions
Colloids (emulsions)
• Are heterogeneous translucent mixtures, e.g.,
cytosol
• Are 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
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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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Chemically Inert and Chemically Reactive Elements
• 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)
• Chemically Inert Elements
 Stable and unreactive
 Outermost energy level fully occupied or contains eight electrons
• 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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(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
(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
Chemical Bonds
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
• Ionic compounds form crystals instead of individual
molecules (Example: NaCl or Sodium chloride)
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Ionic Bond Between Sodium And Chlorine
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
Salt Crystals
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
Reacting atoms
Resulting molecules
+
Hydrogen
atoms
or
Carbon
atom
Molecule of
methane gas (CH4)
(a) Formation of four single covalent bonds:
Structural
formula
shows
single
bonds.
Carbon shares four electron pairs with four hydrogen atoms.
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Double Covalent Bonds
Reacting atoms
Resulting molecules
+
Oxygen
atom
or
Oxygen
atom
Molecule of
oxygen gas (O2)
Structural
formula
shows
double
bond.
(b) Formation of a double covalent bond: Two
oxygen atoms share two electron pairs.
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Figure 2.7b
Triple Covalent Bonds
Reacting atoms
Resulting molecules
+
Nitrogen
atom
or
Nitrogen
atom
Molecule of
nitrogen gas (N2)
Structural
formula
shows
triple
bond.
(c) Formation of a triple covalent bond: Two
nitrogen atoms share three electron pairs.
Figure 2.7c
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Non-polar Covalent Bonds
• Sharing of electrons may be equal or unequal
• 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
 Equal sharing of electrons produces electrically balanced nonpolar
molecules
Example: CO2
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Polar Covalent Bonds
 Unequal sharing by atoms with different electronattracting abilities produces polar molecules
Example: Water
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Figure 2.9
Hydrogen Bonds
• A hydrogen bond forms when a hydrogen
atom covalently bonded to one
electronegative atom is also attracted to
another electronegative atom
 In living cells, the electronegative partners
of hydrogen are usually oxygen or
nitrogen atoms
• Hydrogen bonds are common between dipoles
such as water
• Hydrogen bonds also act as intramolecular
bonds, holding a large molecule in a threedimensional shape
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Hydrogen Bonding Between Water Molecules
+
–
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
Examples:
H + H  H2 (hydrogen gas)
4H + C  CH4 (methane)
(reactants)
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(product)
Patterns of Chemical Reactions
(a) Synthesis reactions
• Synthesis (combination)
reactions
• Decomposition reactions
• Exchange reactions
Synthesis Reactions
• A + B  AB
• Always involve bond
formation
• Anabolic
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
Figure 2.11a
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Decomposition Reactions
• 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
Exchange Reactions
(c) Exchange reactions
• AB + C  AC + B
• Also called
displacement
reactions
• Bonds are both
made and broken
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
Adenosine diphosphate (ADP)
phosphate
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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 a chemical 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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