CH.2 Flashcards
Matter
—anything that has mass and
occupies space
3 states of matter
solid, liquid, gas
Energy
capacity to do work
kinetic
energy in action
potential
stored 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
Travels in waves (e.g., visible light, ultraviolet
light, and x-rays)
Elements
– Matter is composed of elements
– Elements cannot be broken into simpler
substances by ordinary chemical methods
– Each has unique properties
• Atoms
Unique building blocks for each element – Give each element its physical & chemical properties – Smallest particles of an element with properties of that element
Atomic symbol
One- or two-letter chemical shorthand for
each element
Four elements make up 96.1% of body mass
Element Carbon Hydrogen Oxygen Nitrogen
9 elements make up 3.9% of body mass
Element Calcium Phosphorus Potassium Sulfur Sodium Chlorine Magnesium Iodine Iron
11 elements make up < 0.01% of body mass
Element Chromium Copper Fluorine Manganese Silicon Zinc
Atoms
composed of subatomic particles – Protons, neutrons, electrons • Protons and neutrons found in nucleus • Electrons orbit nucleus in an electron cloud
Nucleus of atoms
Almost entire mass of the atom • Neutrons • Carry no charge • Mass = 1 atomic mass unit (amu) • Protons • Carry positive charge • Mass = 1 amu
Electrons in orbitals within electron cloud
• Electrons in orbitals within electron cloud
– Carry negative charge
– 1/2000 the mass of a proton (0 amu)
– Number of protons and electrons always
equal
Different elements contain different
numbers of subatomic particles
Hydrogen has 1 proton, 0 neutrons, and 1
electron
– Lithium has 3 protons, 4 neutrons, and 3
electrons
Atomic number
Number of protons in
nucleus
– Written as subscript to left of atomic symbol Ex. 3Li
Mass number
Total number of protons and neutrons in nucleus • Total mass of atom – Written as superscript to left of atomic symbol • Ex. 7Li
• Isotopes
– Structural variations of atoms
– Differ in the number of neutrons they contain
– Atomic numbers same; mass numbers
different
Atomic weight
Average of mass numbers (relative weights)
of all isotopes of an atom
Radioisotopes - Heavy isotopes decompose to more stable forms
Spontaneous decay called radioactivity
– Similar to tiny explosion
– Can transform to different element
Can be detected with scanners
Radioisotopes
Valuable tools for biological research and medicine – Share same chemistry as their stable isotopes – Most used for diagnosis • All damage living tissue – Some used to destroy localized cancers – Radon from uranium decay causes lung cancer
Molecule
• Two or more atoms bonded together (e.g., H2 or
C6H12O6
)
• Smallest particle of a compound with specific
characteristics of the compound
Compound
Two or more different kinds of atoms bonded
together (e.g., C6H12O6
, but not H2)
Three types of mixtures
– Solutions
– Colloids
– Suspensions
Solvent
– Substance present in greatest amount
– Usually a liquid; usually water
Solute
– Present in smaller amounts
• Ex. If glucose is dissolved in blood, glucose is
solute; blood is solvent
Colloids
Heterogeneous mixtures, e.g., cytosol
– Large solute particles do not settle out
– Some undergo sol-gel transformations
• e.g., cytosol during cell division
Suspensions
Heterogeneous mixtures, e.g., blood
– Large, visible solutes settle out
• Mixtures
No chemical bonding between components
– Can be separated by physical means, such as
straining or filtering
– Heterogeneous or homogeneous
Compounds
Chemical bonding between components
– Can be separated only by breaking bonds
– All are homogeneous
• Chemical bonds
are energy relationships
between electrons of reacting atoms
• Electrons in valence shell (outermost electron
shell)
Have most potential energy
– Are chemically reactive electrons
Chemically Reactive Elements
• Valence shell not full
• Tend to gain, lose, or share electrons
(form bonds) with other atoms to achieve
stability
Types of Chemical Bonds
– Ionic bonds
– Covalent bonds
– Hydrogen bonds
Ions
Atom gains or loses electrons and becomes charged
• # Protons ≠ # Electrons
• Transfer of valence shell electrons from one
atom to another forms ions
– One becomes an anion (negative charge)
• Atom that gained one or more electrons
– One becomes a cation (positive charge)
• Atom that lost one or more electrons
• Attraction of opposite charges results in an ionic
bond
Ionic Compounds
• Most ionic compounds are salts
– When dry salts form crystals instead of
individual molecules
– Example is NaCl (sodium chloride)
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
Nonpolar Covalent Bonds
Electrons shared equally
• Produces electrically balanced, nonpolar
molecules such as CO2
Polar Covalent Bonds
• Unequal sharing of electrons produces polar (AKA dipole) molecules such as H2O – Atoms in bond have different electronattracting abilities • Small atoms with six or seven valence shell electrons are electronegative, e.g., oxygen – Strong electron-attracting ability • Most atoms with one or two valence shell electrons are electropositive, e.g., sodium
Hydrogen Bonds
Attractive force between electropositive
hydrogen of one molecule and an
electronegative atom of another molecule
– Not true bond
– Common between dipoles such as water
– Also act as intramolecular bonds, holding a
large molecule in a three-dimensional shape
Chemical Reactions
• Occur when chemical bonds are formed,
rearranged, or broken
• Represented as chemical equations using
molecular formulas
– Subscript indicates atoms joined by bonds
– Prefix denotes number of unjoined atoms or
molecules
• Chemical equations contain
– Reactants
• Number and kind of reacting substances
– Chemical composition of the product(s)
– Relative proportion of each reactant and product in
balanced equations
Patterns of Chemical Reactions
- Synthesis (combination) reactions
- Decomposition reactions
- Exchange reactions
Decomposition Reactions
AB A + B – Molecule is broken down into smaller molecules or its constituent atoms • Reverse of synthesis reactions – Involve breaking of bonds – Catabolic
Exchange Reactions
AB + C AC + B
– Also called displacement reactions
– Involve both synthesis and decomposition
– Bonds are both made and broke
Oxidation-Reduction (Redox) Reactions
Are decomposition reactions
– Reactions in which food fuels are broken down for
energy
• Are also exchange reactions because electrons
are exchanged between reactants
– Electron donors lose electrons and are oxidized
– Electron acceptors receive electrons and become
reduced
• C6H12O6 + 6O2 6CO2 + 6H2O + ATP
• Glucose is oxidized; oxygen molecule is reduced
Exergonic reactions
net release of energy
• Products have less potential energy than reactants
• Catabolic and oxidative reactions
Endergonic reactions
net absorption of energy • Products have more potential energy than reactants • Anabolic reactions
Rate of Chemical Reactions
Affected by – Temperature Rate – Concentration of reactant Rate – Particle size Rate – Catalysts: Rate without being chemically changed or part of product • Enzymes are biological catalysts
Biochemistry
Study of chemical composition and
reactions of living matter
• All chemicals either organic or inorganic
• Inorganic compounds
Water, salts, and many acids and bases
• Do not contain carbon
Organic compounds
Carbohydrates, fats, proteins, and nucleic
acids
• Contain carbon, usually large, and are
covalently bonded
Polar solvent properties of water
– Dissolves and dissociates ionic substances
– Forms hydration layers around large charged
molecules, e.g., proteins (colloid formation)
– Body’s major transport medium
Salts
Ionic compounds that dissociate into ions in
water
– Ions (electrolytes) conduct electrical currents in
solution
– Ions play specialized roles in body functions (e.g.,
sodium, potassium, calcium, and iron)
– Ionic balance vital for homeostasis
• Contain cations other than H+ and anions other
than OH–
• Common salts in body
– NaCl, CaCO3
, KCl, calcium phosphates
Acids
s are proton donors – Release H+ (a bare proton) in solution – HCl H+ + Cl– is a electrolyte
Bases
proton acceptors – Take up H+ from solution • NaOH Na+ + OH– – OH– accepts an available proton (H+ ) – OH– + H+ H2O electrolyte
Important acids
– HCl, HC2H3O2
(HAc), and H2CO
• Important bases
– Bicarbonate ion (HCO3–) and ammonia (NH3
pH: Acid-base Concentration
– Relative free [H+ ] of a solution measured on pH scale – As free [H+ ] increases, acidity increases • [OH– ] decreases as [H+ ] increases • pH decreases – As free [H+ ] decreases alkalinity increases • [OH– ] increases as [H+ ] decreases • pH increases
pH: Acid-base Concentration
Acidic solutions [H+ ], pH – Acidic pH: 0–6.99 • Neutral solutions – Equal numbers of H+ and OH– – All neutral solutions are pH 7 – Pure water is pH neutral • pH of pure water = pH 7: [H+ ] = 10–7 m • Alkaline (basic) solutions [H+ ], pH – Alkaline pH: 7.01–14
Buffers
Acidity reflects only free H+
in solution
– Not those bound to anions
• Buffers resist abrupt and large swings in pH
– Release hydrogen ions if pH rises
– Bind hydrogen ions if pH falls
• Convert strong (completely dissociated) acids or bases
into weak (slightly dissociated) ones
• Carbonic acid-bicarbonate system (important buffer
system of blood):
Organic Compounds
Molecules that contain carbon
– Except CO2 and CO, which are considered
inorganic
– Carbon is electroneutral
• Shares electrons; never gains or loses them
• Forms four covalent bonds with other elements
• Unique to living systems
• Carbohydrates, lipids, proteins, and
nucleic acids
Many are polymers
– Chains of similar units called monomers
(building blocks)
• Synthesized by dehydration synthesis
• Broken down by hydrolysis reactions
Carbohydrates
• Sugars and starches • Polymers • Contain C, H, and O [(CH20)n ] • Three classes – Monosaccharides – one sugar – Disaccharides – two sugars – Polysaccharides – many sugars Functions of carbohydrates – Major source of cellular fuel (e.g., glucose) – Structural molecules (e.g., ribose sugar in RNA)
Monosaccharides
Simple sugars containing three to seven C atoms • (CH20)n – general formula; n = # C atoms • Monomers of carbohydrates • Important monosaccharides – Pentose sugars • Ribose and deoxyribose – Hexose sugars • Glucose (blood sugar)
Disaccharides
Double sugars
• Too large to pass through cell membranes
• Important disaccharides
– Sucrose, maltose, lactose
Polysaccharides
• Polymers of monosaccharides
• Important polysaccharides
– Starch and glycogen
• Not very soluble
Lipids
Contain C, H, O (less than in carbohydrates), and sometimes P • Insoluble in water • Main types: – Neutral fats or triglycerides – Phospholipids – Steroids – Eicosanoids
Neutral Fats or Triglycerides
• Called fats when solid and oils when liquid
• Composed of three fatty acids bonded to A
glycerol molecule
• Main functions
– Energy storage
– Insulation
– Protection
Saturation of Fatty Acids
Saturated fatty acids – Single covalent bonds between C atoms • Maximum number of H atoms – Solid animal fats, e.g., butter • Unsaturated fatty acids – One or more double bonds between C atoms • Reduced number of H atoms – Plant oils, e.g., olive oil – “Heart healthy” • Trans fats – modified oils – unhealthy • Omega-3 fatty acids – “heart healthy”
Phospholipids
Modified triglycerides:
– Glycerol + two fatty acids and A phosphorus
(P) - containing group
• “Head” and “tail” regions have different
properties
• Important in cell membrane structure
Steroids
Steroids—interlocking four-ring structure
• Cholesterol, vitamin D, steroid hormones,
and bile salts
• Most important steroid
– Cholesterol
• Important in cell membranes, vitamin D synthesis,
steroid hormones, and bile salts
Eicosanoids
• Many different ones
• Derived from a fatty acid (arachidonic
acid) in cell membranes
• Most important eicosanoid
– Prostaglandins
• Role in blood clotting, control of blood pressure,
inflammation, and labor contractions
Proteins
Contain C, H, O, N, and sometimes S and P
• Proteins are polymers
• Amino acids (20 types) are the monomers in
proteins
– Joined by covalent bonds called peptide bonds
– Contain amine group and acid group
– Can act as either acid or base
– All identical except for “R group” (in green on figure)
Fibrous (structural) proteins
Strandlike, water-insoluble, and stable
– Most have tertiary or quaternary structure (3-D)
– Provide mechanical support and tensile
strength
– Examples: keratin, elastin, collagen (single
most abundant protein in body), and certain
contractile fibers
Globular (functional) proteins
Compact, spherical, water-soluble and
sensitive to environmental changes
– Tertiary or quaternary structure (3-D)
– Specific functional regions (active sites)
– Examples: antibodies, hormones, molecular
chaperones, and enzymes
Protein Denaturation
Globular proteins unfold and lose functional,
3-D shape
• Active sites destroyed
– Can be cause by decreased pH or increased
temperature
• Usually reversible if normal conditions
restored
• Irreversible if changes extreme
– e.g., cooking an egg
Molecular Chaperones
Globular proteins • Ensure quick, accurate folding and association of other proteins • Prevent incorrect folding • Assist translocation of proteins and ions across membranes • Promote breakdown of damaged or denatured proteins • Help trigger the immune response
• Stress proteins
– Molecular chaperones produced in response
to stressful stimuli, e.g., O2 deprivation
– Important to cell function during stress
– Can delay aging by patching up damaged
proteins and refolding them
Enzymes
Globular proteins that act as biological
catalysts
• Regulate and increase speed of chemical
reactions
– Lower the activation energy, increase the
speed of a reaction (millions of reactions per
minute!)
Characteristics of Enzymes
• Some functional enzymes (holoenzymes) consist of two parts – Apoenzyme (protein portion) – Cofactor (metal ion) or coenzyme (organic molecule often a vitamin) • Enzymes are specific – Act on specific substrate • Usually end in -ase • Often named for the reaction they catalyze – Hydrolases, oxidases
Nucleic Acids
Deoxyribonucleic acid (DNA) and ribonucleic acid (RNA) – Largest molecules in the body • Contain C, O, H, N, and P • Polymers – Monomer = nucleotide • Composed of nitrogen base, a pentose sugar, and a phosphate group
Deoxyribonucleic Acid (DNA)
Utilizes four nitrogen bases:
– Purines: Adenine (A), Guanine (G)
– Pyrimidines: Cytosine (C), and Thymine (T)
– Base-pair rule – each base pairs with its
complementary base
• A always pairs with T; G always pairs with C
• Double-stranded helical molecule (double helix)
in the cell nucleus
• Pentose sugar is deoxyribose
• Provides instructions for protein synthesis
• Replicates before cell division ensuring genetic
continuity
Ribonucleic Acid (RNA)
Four bases:
– Adenine (A), Guanine (G), Cytosine (C), and
Uracil (U)
• Pentose sugar is ribose
• Single-stranded molecule mostly active
outside the nucleus
• Three varieties of RNA carry out the DNA
orders for protein synthesis
– Messenger RNA (mRNA), transfer RNA
(tRNA), and ribosomal RNA (rRNA)
Adenosine Triphosphate (ATP)
• Chemical energy in glucose captured in
this important molecule
• Directly powers chemical reactions in cells
• Energy form immediately useable by all
body cells
• Structure of ATP
– Adenine-containing RNA nucleotide with two
additional phosphate group
• Phosphorylation
Terminal phosphates are enzymatically
transferred to and energize other molecules
– Such “primed” molecules perform cellular
work (life processes) using the phosphate
bond energy
