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Assign: Periodic Table Basics - Review
1
Basic Chemistry
• Matter, Mass, and Weight
– Matter: anything that occupies space and has
mass
– Mass: the amount of matter in an object
– Weight: the gravitational force acting on an
object of a given mass
• Composition of Matter: Elements and Atoms
– Element: the simplest type of matter with
unique chemical properties
• composed of atoms of only one kind
– Atom: smallest particle of an element that has
chemical characteristics of that element
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Atomic Structure
• Atoms: composed of
subatomic particles
– Neutrons: no electrical
charge
– Protons: one positive
charge
– Electrons: one negative
charge
• Nucleus: formed by
protons and neutrons
• Most of the volume of an
atom is occupied by
electrons
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• Elements are
arranged by atomic
number
• 24 elements have a
biological role
• 4 elements = 96.2%
of body wt.
•7 elements = 3.7% of
body wt.
• Trace elements in
minute amounts
(0.1%)
The four most important elements (constituting 96.2% of
body weight) include:
Oxygen (65%)
Carbon (18.5%)
Hydrogen (9.5%)
Nitrogen (3.2%)
Less important elements (constituting 3.7% of body
weight) include:
Calcium (1.5%)
Phosphate (1.0%)
Potassium (0.4%)
Sulphur (0.3%)
Sodium (0.2%)
Chlorine (0.2%)
Magnesium (0.1%)
The remaining 0.1% belongs to the family of trace
elements:
iron, manganese, zinc, copper, iodine, cobalt,
molybdenum, selenium, chromium, silicon, fluorine,
vanadium, nickel, arsenic, and tin.
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Atomic Number and Mass Number
• Atomic Number
– Equal to number of protons
in each atom
– Equal to the number of
electrons
• Mass (Atomic) Number
– Number of protons plus
number of neutrons
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Identification of Elements
6 2.2
Figure
Isotopes
•
•
Isotopes: two or more forms of the same element with same number of
protons and electrons but different neutron number
– For example; there are three types of hydrogen; see diagram below
– Denoted by using symbol of element preceded by mass number as 1H,
2H, 3H
Radioactive isotopes
– Forms of atoms that emit radioactivity such as gamma rays, which can then be
measured
– Used clinically and in research
– Examples of uses
• Tracking hormone uptake
• Treating cancer
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Ions and Ionization
• Ions - an atom that carries a charge due to an
unequal number of protons and electrons
• Ionization = transfer
of electrons from one
atom to another
( stability of valence
shell)
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Anions and Cations
• Anion
– atom that gained electrons (net negative charge)
• Cation
– atom that lost an electron (net positive charge)
• Ions with opposite charges are attracted to each
other
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Electrolytes
• Salts that ionize in water to form body fluids
– capable of conducting electricity
• Electrolyte importance
– chemical reactivity
– osmotic effects (influence water movement)
– electrical effects on nerve and muscle tissue
• Imbalances cause muscle cramps, brittle
bones, coma and death
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7 Major Electrolytes In the human body
•
•
•
Sodium (Na+)
Chloride (Cl-)
Potassium (K+)
•
•
•
•
Magnesium(Mg++)
Calcium (Ca++)
Phosphate (HPO4–)
Bicarbonate (HCO3-)
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Molecules and Chemical Bonds
• Molecules
– two or more atoms covalently bonded
– Example: a hydrogen molecule (H2)
• Compounds
– two or more atoms of different elements covalently bonded Example:
water (H2O)
• Molecular formula
– elements and how many atoms of each
– E.g., H2O - 2 hydrogens +1 oxygen
– E.g. 2 H2O = 2 molecules of water
• Structural formula
– location of each atom
– structural isomers revealed
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Molecular Weight
• MW of compound = sum of atomic weights of
all the atoms in the molecule
• Calculate: MW of glucose (C6H12O6)
6 C atoms x 12 amu each = 72 amu
12 H atoms x 1 amu each = 12 amu
6 O atoms x 16 amu each = 96 amu
Molecular weight (MW) = 180 amu
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Chemical Bonds
• Electron shells, or energy levels,
surround the nucleus of an atom
• Bonds are formed using the electrons in
the outermost energy level
• ONLY electrons are used in forming
chemical bonds
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Chemically Inert Elements
• Inert elements have their outermost energy
level fully occupied by electrons;
nonreactive
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Chemically Reactive Elements
• Reactive elements
do not have their
outermost energy
level fully occupied
by electrons
• Are capable of
sharing, gaining, or
losing electrons
and thus forming
bonds
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Chemical Bonds: Ionic Bonds
• Ionic bonds form between atoms by the
transfer of one or more electrons
– Loss of electron(s) by one atom results in the gain
of the electron(s) by another atom
• Ionic compounds form crystals instead of
individual molecules
• Example: NaCl (sodium chloride)
• Ions are charged atoms resulting from the gain
or loss of electrons
– Cations have lost one or more electrons
• Indicated by a positive charge; e.g., Na+, K+
– Anions have gained one or more electrons
• Indicated by a negative charge; e.g., Cl-
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Formation of an Ionic Bond
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Covalent Bonding
• Atoms share one or
more pairs of electrons
– Single covalent: two
atoms share one pair of
electrons
– Nonpolar covalent:
Electrons shared equally
because nuclei attract the
electrons equally
– Polar covalent: Electrons
not shared equally
because one nucleus
attracts the electrons
more than the other one
does
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Single Covalent Bonds
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Double Covalent Bonds
Assign: Covalent & Ionic
Bonding Assignment
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Polar and Nonpolar Molecules
• Electrons shared equally between
atoms produce nonpolar molecules
• Unequal sharing of electrons produces
polar molecules
• Atoms with six or seven valence shell
electrons are electronegative
• Atoms with one or two valence shell
electrons are electropositive
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Nonpolar /Polar Covalent
Bonds
electrons
shared
equally
electrons shared
unequally
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Hydrogen Bonds
• An intermolecular force
– A bond that forms between molecules rather than
within a molecule
• Weakest bond = no sharing of electrons
• Attraction between polar molecules
– Positive hydrogen atoms attracted to negative oxygen
atoms in a 2nd molecule
• Physiological importance
– play an important role in determining the shape of
complex molecules:
• proteins, nucleic acids
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Hydrogen Bonds
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Solvency
• Solvency - ability to dissolve other chemicals
– Hydrophilic (charged substances) dissolve
easily in water
– Hydrophobic (neutral substances) do not easily
dissolve in water
• Water = universal solvent
– More solutes are soluble in water than any other
solvent
– metabolic reactions and transport of substances
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Sec 2-2: Properties of Water – page 40
1. Universal solvent
– More solutes are soluble in water than in any other
solvent
– Due mainly to its polarity
2. Chemical reactivity
– Water participates in chemical reactions
•
•
•
water ionizes into H+ and OHwater ionizes other chemicals (acids and salts)
water is involved in hydrolysis (digestion) and
dehydration synthesis (formation) reactions
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Properties of Water: As a Solvent
• Polar water molecules overpower the ionic bond in
Na+Cl– forming hydration spheres around each ion
– water molecules: negative pole faces Na+, positive pole
faces Cl28
Properties of Water: Thermal
Stability of Water
3. Water used to transport solutes and heat
– 90-92% of plasma is water
– All solutes and heat are transported in blood
• Allows heat generated in one region to be disbursed over
entire body
4. Water stabilizes internal temperature
– Has high heat capacity
• hydrogen bonds inhibit temperature increases by inhibiting
molecular motion
– water can absorb large amounts of heat without large changes in
temperature
– Has high heat of vaporization (see: http://health.howstuffworks.com/sweat2.htm)
• effective coolant
• 1 ml of sweat removes 500 calories of heat
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Electrolytes and Nonelectrolytes
• Electrolytes: solutions made by the
dissociation of cations (+) and anions (-)
in water
– Have the capacity to conduct an electric
current
– Currents can be detected by electrodes
• Nonelectrolytes: solutions made by
molecules that dissolve in water, but do
not dissociate; do not conduct
electricity
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Acids and Bases; Salts and Buffers page42
• Acid: a proton (H+) donor or any substance
that releases hydrogen ions: HCl  H+ + Cl –
• Base: a proton acceptor or any substance
that binds to or accepts hydrogen ions
– NaOH  Na+ + OH–
• Salt: a compound consisting of a cation other
than a hydrogen ion and an anion other than a
hydroxide ion. Example: NaCl, KCl
• Buffer: a substance which prevents wide
fluctuations in pH
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Acid-Base Concentration (pH)
• Acidic solutions have higher H+
concentration and therefore a lower pH
– pH from 0-7
• Alkaline solutions have lower H+
concentration and therefore a higher pH
– pH from 7-14
• Neutral solutions have equal H+ and OH–
concentrations
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The pH Scale
• Refers to the hydrogen ion
concentration in a
solution
– Neutral: pH of 7 or
equal hydrogen and
hydroxide ions
– Acidic: a greater
concentration of
hydrogen ions
– Alkaline or basic: a
greater
concentration of
hydroxide ions 33
Sec 2-4:
Chemical Reactions page 48
• Atoms, ions, molecules or compounds
interact to form or break chemical bonds
– Reactants: substances that enter into a
chemical reaction.
– Products: substances that result from the
reaction
• Chemical bonds are made (synthesis;
anabolism) and broken (decomposition;
catabolism) during chemical reactions
• Chemical equations contain:
– Number and type of reacting substances, and
products produced
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Synthetic Reactions
• Two or more reactants chemically combine to form a
new and larger product. Anabolism.
– Chemical bonds made; energy stored in the bonds.
– Responsible for growth, maintenance and repair
– Dehydration synthesis: synthesis reaction where water is a
product
– Produce chemicals characteristic of life: carbohydrates,
proteins, lipids, and nucleic acids
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Decomposition Reactions
• A large reactant is broken down to form smaller products.
Catabolism.
– Chemical bonds broken; energy released.
– Hydrolysis: water is split into two parts that contribute to the
formation of the products
– Example: the breakdown of ATP to form ADP and inorganic
phosphate with a concomitant release of free energy
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Reversible Reactions
• Chemical reactions in which the
reaction can proceed either from
reactants to products or from products
to reactants.
• Equilibrium: rate of product formation is
equal to rate of reactant formation
• Example: CO2 and H+ formation in
plasma
CO2 + H2O  H2CO3  H+ + HCO337
Oxidation-Reduction (Redox)
Reactions
• Oxidation
– molecule loses electrons and releases energy
– oxygen is often the electron acceptor
• Reduction
– molecule gains electrons and energy
• Oxidation-reduction (redox) reactions
– Electrons are often transferred as hydrogen atoms
Citric acid
α-ketoglutaric acid
-2H
NAD
+2H
NADH + H
– NAD and FAD are commonly used to accept hydrogens from cmpds
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undergoing oxidation (like citric acid in the above example)
Organic Chemistry: Biochemicals
• Carbohydrates: composed of carbon, hydrogen, oxygen.
– Divided into monosaccharides, disaccharides, polysaccharides
– Example: glucose
– Energy sources and structure
• Lipids: composed mostly of carbon, hydrogen, oxygen.
– Relatively insoluble in water.
– Example: anabolic steroids
– Functions: protection, insulation, physiological regulation, component
of cell membranes, energy source
• Proteins: composed of carbon, hydrogen, oxygen, nitrogen
– Example: insulin
– Functions: regulate processes, aid transport, protection, muscle
contraction, structure, energy
• Nucleic Acids: composed of carbon, hydrogen, oxygen,
nitrogen, and phosphorus.
– Examples: ATP, DNA, RNA
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Sec 2-3: Organic Molecules
and Carbon page 44
• Carbon has only 4 valence electrons
– bonds readily to gain 4 more valence electrons
• Forms long chains, branched molecules
and rings
– serves as the backbone for organic molecules
• Carries a variety of functional groups
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Functional Groups
• Atoms attached to
carbon backbone
• Determines chemical
properties of the cmpd
to which it is attached
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Carbohydrates: Monosaccharides
•
•
Simple sugars
General formula is CH2O
• Six-carbon sugars (hexoses) like glucose, fructose, and galactose
are important in the diet as energy sources (C6H12O6)
– structural isomers: same molecular but different structural formula
• Five-carbon sugars (pentoses) are components of ATP, DNA and
RNA; ribose and deoxyribose
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Carbohydrates: Disaccharides
• Two simple sugars bound together by dehydration
synthesis (loss of water with formation of new cmpd
• Examples: sucrose (Glu + Fru), lactose (Glu + Galac),
maltose (Glu + Glu)
• Are isomers: C12H22O11: Why not C12H24O12?
• Digested through hydrolysis reactions
– Note the reversible reaction below
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Carbohydrates: Polysaccharides
• Chains of glucose subunits
Polymers
• Glycogen formed by
animals; energy storage
– Liver synthesizes after a
meal and breaks it down
between meals
• Starch and cellulose
formed by plants
– Starch in food is used as a
source of monosaccharides
– Cellulose in food acts as fiber
(bulk) in the diet
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Carbohydrate Functions
• All digested carbohydrates converted to glucose and oxidized
to make ATP
• Conjugated carbohydrate = bound to lipid or protein
– glycolipids
• external surface of cell membrane
– glycoproteins
• external surface of cell membrane
• mucus of respiratory and digestive tracts
– proteoglycans
• gels that hold cells and tissues together
• joint lubrication
• rubbery texture of cartilage
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Lipids: Fats
• Contain C, H, and O, but the proportion of oxygen in
lipids is less than in carbohydrates
• Hydrophobic organic molecules
• Examples: Triglycerides, Phospholipids, Steroids,
Eicosanoids
• Ingested and broken down by hydrolysis
– Triglycerides: composed of glycerol and 3 fatty acids
• Fatty acids may be saturated or unsaturated
• Saturated (no double bonds between carbons)
• Unsaturated (one or more double bonds between
carbon atoms)
– Functions: protection, insulation, energy source
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Lipids: Triglycerides
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Lipids: Phospholipids
• Phospholipids – modified triglycerides with two fatty acid
groups and a phosphorus group
• Polar (hydrophilic) at one end; nonpolar (hydrophobic) at
the other.
– Function: important structural component of cell membranes
48
Lipids: Steroids
• Steroids – flat molecules with four interlocking hydro-
carbon rings
• All are derived from cholesterol
• Cholesterol
– important
component of cell
membranes
– produced only in
animal liver cells
• naturally produced
by our body
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Representative Lipids Found
in the Body
• Neutral fats – found in subcutaneous tissue and
around organs
• Phospholipids – chief component of cell membranes
• Steroids – cholesterol, bile salts, vitamin D, sex
hormones, and adrenal cortical hormones
• Fat-soluble vitamins – vitamins A, D, E, and K
• Eicosanoids – prostaglandins, leukotriens, and
thromboxanes
• Lipoproteins – transport fatty acids and cholesterol in
the bloodstream
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Proteins
Macromolecules composed of combinations of 20
types of amino acids bound together by peptide bonds
• Contain C, H, O, N
• Amino acids: building
blocks of protein
– Contain an amino group
(NH2), carboxyl (COOH)
group, and a radical group
(R-)
• Peptide bonds: covalent
bonds formed between
amino acids during
protein synthesis
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Levels of Protein Structure
• Primary: sequence of
amino acids in the
polypeptide chain
• Secondary: folding and
bending of chain caused by
hydrogen bonding
• Tertiary: formation of
helices or of pleated
sheets; caused in part by
S-S bonds between amino
acids
• Quaternary: two or more
proteins associate as a
functional unit
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Fibrous and Globular Proteins
• Fibrous proteins
– Extended and strand-like proteins
– Examples: keratin, elastin, collagen, and
certain contractile fibers
• Globular proteins
– Compact, spherical proteins with tertiary
and quaternary structures
– Examples: antibodies, hormones, and
enzymes
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Types of Proteins (8)
1.
2.
3.
4.
5.
6.
7.
8.
Structural - collagen, keratin
Transport - Hb,  and  globulins
Contractile - actin and myosin in muscle
Regulatory - hormones
Immunologic - antibodies (IgG, IgA)
Clotting - thrombin and fibrin
Osmotic - albumin in plasma
Catalytic - enzymes
54
Protein Denuaturation
• Reversible
unfolding of
proteins due
to drops in
pH and/or
increased
temperature
55
Protein Denaturation
• Irreversibly denatured proteins cannot refold
and are formed by extreme pH or
temperature changes
56
Nucleic Acids
• Composed of carbon, oxygen, hydrogen,
nitrogen, and phosphorus
• Their structural unit, the nucleotide, is
composed of N-containing base, a pentose
sugar, and a phosphate group
• Five nitrogen bases contribute to nucleotide
structure – adenine (A), guanine (G), cytosine
(C), thymine (T), and uracil (U)
• Two major classes – DNA and RNA
– Informational molecules
• Other important nucleotides:
– ATP, cAMP
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DNA: Deoxyribonucleic acid
• 100 million to 1 billion
nucleotides long
• Genetic material of cells
copied from one
generation to next
• Composed of 2 strands
of nucleotides
– Each nucleotide contains
one of the organic bases
of adenine or guanine
(which are purines) and
thymine or cystosine
(which are pyrimidines).
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RNA: Ribonucleic acid
• Similar to a single strand of DNA
– Four different nucleotides make up organic
bases except thymine is replaced with
uracil (pyrimidine)
• Responsible for interpreting the code
within DNA into the primary structure of
proteins.
59
Adenosine Triphosphate
(ATP)
• Energy currency of the body
• Provides energy for other chemical reactions as
anabolism or drive cell processes as muscle
contraction
• All energy-requiring chemical reactions stop when
there is inadequate ATP
60
How ATP Drives Cellular Work
61
Turnover Rates (TR)of
Different Cell Types
• TR refers to the average time between synthesis and recycling of
the compound
• Liver
–
–
–
–
Total protein: 5-6 days
Enzymes: 1 hr to several days
Glycogen: 1-2 days
Cholesterol: 5-7 days
• Muscle
– Total protein: 30 days
– Glycogen: 12-24 hours
• Neurons
– Phospholipids: 200 days
– Cholesterol: 100+ days
• Fat cell
– Triglycerides: 15-20 days
62
Example Reaction: Getting rid
of Carbon Dioxide
• In the blood
CO2 + H20  H2CO3 (carbonic acid)
• In the lungs
H2CO3  CO2 + H2O
Released as you breathe
Energy in reactions
Energy-Absorbing Reaction
Energy-Releasing Reaction
Activation
energy
Products
Activation energy
Reactants
Reactants
Products
Activation Energy
• The energy that is needed to get
a reaction started
Enzymes
• Some chemical reactions are too
slow or have activation energies
that are too high to make them
practical for living tissue.
• These chemical reactions are
made possible by CATALYSTS.
Catalyst
• Substance that speeds up the
rate of chemical reactions
• Work by lowering a reactions
activation energy
Enzyme
•
•
•
•
•
BIOLOGICAL CATALYSTS
Speed up reactions in cells
Very specific
Named for the reaction it catylzes
Enzyme names always end in -ase
Reaction pathway
without enzyme
Reactants
Reaction pathway
with enzyme
Activation energy
without enzyme
Activation
energy
with enzyme
Products
Substrates
• The reactants of enzyme
catalyzed reactions
• The active site of the enzyme
and the substrate have
complementary shapes
• Fit like a lock and key
Enzyme Action
Enzyme – substrate complex
Enzyme
(hexokinase)
Glucose
ADP
Substrates
Products
ATP
Glucose-6phosphate
Products
are released
Active site
Enzyme-substrate
complex
Substrates
are converted
into products
Substrates
bind to
enzyme
Regulation of Enzyme Activity
• Enzymes are affected by any
variable that affects chemical
reactions
1. pH
2. Temperature
3. Concentration
of enzyme