Chemical Foundationsfor Cells
Chapter 2
Elements
•Fundamental forms of matter
•Can’t be broken apart by normalmeans
•92 occur naturally on Earth
How is matter organized
Most Common Elements inLiving Organisms
Oxygen
Hydrogen
Carbon
Nitrogen
What Are Atoms?
•Smallest particles that retain propertiesof an element
•Made up of subatomic particles:
–Protons (+)
–Electrons (-)
–Neutrons (no charge)
Atoms
Fig. 2.3, p. 22
HYDROGEN
HELIUM
electron
proton
neutron
Hydrogen and Helium Atoms
Atomic Number
•Number of protons
•All atoms of an element have the sameatomic number
•Atomic number of hydrogen = 1
•Atomic number of carbon = 6
Mass Number
Number of protons
+
Number of neutrons
Isotopes vary in mass number
Atomic Number
Isotopes
•Atoms of an element with differentnumbers of neutrons (different massnumbers)
•Carbon 12 has 6 protons, 6 neutrons
•Carbon 14 has 6 protons, 8 neutrons
Radioisotopes
•Have an unstable nucleus thatemits energy and particles
•Radioactive decay transformsradioisotope into a different element
•Decay occurs at a fixed rate
isotopes
Radioisotopes as Tracers
•Tracer is substance with aradioisotope attached to it
•Emissions from the tracer can bedetected with special devices
•Following movement of tracers isuseful in many areas of biology
Thyroid Scan
•Measures health of thyroid by detectingradioactive iodine taken up by thyroidgland
normal thyroid
enlarged
cancerous
Other Uses of Radioisotopes
•Drive artificial pacemakers
•Radiation therapy
Emissions from some radioisotopes candestroy cells. Some radioisotopes are usedto kill small cancers.
What Determines WhetherAtoms Will Interact?
The number and arrangement of theirelectrons
Electrons
•Carry a negative charge
•Repel one another
•Are attracted to protons inthe nucleus
•Move in orbitals - volumesof space that surround thenucleus
Z
X
When all p orbitals are full
y
Electron Vacancies
•Unfilled shells makeatoms likely to react
•Hydrogen, carbon,oxygen, andnitrogen all havevacancies in theirouter shells
CARBON
6p+ , 6e-
NITROGEN
7p+ , 7e-
HYDROGEN
1p+ , 1e-
Chemical Bonds, Molecules,& Compounds
•Bond is union between electronstructures of atoms
•Atoms bond to form molecules
•Molecules may contain atoms of onlyone element - O2
•Molecules of compounds contain morethan one element - H2O
Chemical Bookkeeping
•Use symbols for elements when writingformulas
•Formula for glucose is C6H12O6
– 6 carbons
– 12 hydrogens
– 6 oxygens
Chemical Bookkeeping
•Chemical equation shows reaction
Reactants ---> Products
•Equation for photosynthesis:
6CO2 + 6H2O ---> + C6H12O12 + 6H2O
Important Bonds in BiologicalMolecules
Ionic Bonds
Covalent Bonds
Hydrogen Bonds
Chemical Bonds
Ion Formation
•Atom has equal number ofelectrons and protons - no netcharge
•Atom loses electron(s), becomespositively charged ion
•Atom gains electron(s), becomesnegatively charged ion
Ionic Bonding
•One atom loses electrons,becomes positively charged ion
•Another atom gains theseelectrons, becomes negativelycharged ion
•Charge difference attracts thetwo ions to each other
Formation of NaCl
•Sodium atom (Na)
–Outer shell has one electron
•Chlorine atom (Cl)
–Outer shell has seven electrons
•Na transfers electron to Cl forming Na+and Cl-
•Ions remain together as NaCl
7mm
SODIUM
ATOM
11 p+
11 e-
SODIUM
ION
11 p+
10 e-
electron transfer
CHLORINE
ATOM
17 p+
17 e-
CHLORINE
ION
17 p+
18 e-
Fig. 2.10a, p. 26
Formation of NaCl
Covalent Bonding
Atoms share a pair or pairs ofelectrons to fill outermost shell
•Single covalent bond
•Double covalent bond
•Triple covalent bond
Nonpolar Covalent Bonds
•Atoms share electrons equally
•Nuclei of atoms have samenumber of protons
•Example: Hydrogen gas (H-H)
Polar Covalent Bonds
•Number of protons in nuclei ofparticipating atoms is NOT equal
•Electrons spend more time nearnucleus with most protons
•Water - Electrons more attractedto O nucleus than to H nuclei
Hydrogen Bonding
•Molecule held together by polarcovalent bonds has no NET charge
•However, atoms of the molecule carrydifferent charges
•Atom in one polar covalent moleculecan be attracted to oppositely chargedatom in another such molecule
one
large
molecule
another
large
molecule
a large
molecule
twisted
back
on
itself
Fig. 2.12, p. 27
Examples of Hydrogen Bonds
Properties of Water
Polarity
Temperature-Stabilizing
Cohesive
Solvent
Water Is a PolarCovalent Molecule
•Molecule has no netcharge
•Oxygen end has aslight negative charge
•Hydrogen end has aslight positive charge
O
H
H
O
H
H
O
H
H
+
_
+
+
+
_
+
+
Liquid Water
Hydrophilic & HydrophobicSubstances
•Hydrophilic substances
–Polar
–Hydrogen bond with water
–Example: sugar
•Hydrophobic substances
–Nonpolar
–Repelled by water
–Example: oil
Temperature-StabilizingEffects
•Liquid water can absorb much heatbefore its temperature rises
•Why?
•Much of the added energy disruptshydrogen bonding rather thanincreasing the movement of molecules
Evaporation of Water
•Large energy input can cause individualmolecules of water to break free into air
•As molecules break free, they carryaway some energy (lower temperature)
•Evaporative water loss is used bymammals to lower body temperature
Why Ice Floats
•In ice, hydrogen bonds lock moleculesin a lattice
•Water molecules in lattice are spacedfarther apart then those in liquid water
•Ice is less dense than water
Water Cohesion
•Hydrogen bonding holdsmolecules in liquid watertogether
•Creates surface tension
•Allows water to move ascontinuous columnupward through stems ofplants
Water Is a Good Solvent
•Ions and polar molecules dissolve easilyin water
•When solute dissolves, water moleculescluster around its ions or molecules andkeep them separated
Fig. 2.16, p. 29
Na+
Cl–
–
–
–
–
–
–
–
–
–
–
–
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
Spheres of Hydration
Hydrogen Ions: H+
•Unbound protons
•Have important biological effects
•Form when water ionizes
The pH Scale
•Measures H+ concentration of fluid
•Change of 1 on scale means 10Xchange in H+ concentration
Highest H+ Lowest H+
0---------------------7-------------------14
Acidic Neutral Basic
Acids and Bases
Examples of pH
•Pure water is neutral with pH of 7.0
•Acidic
–Stomach acid: pH 1.0 - 3.0
–Lemon juice: pH 2.3
•Basic
–Seawater: pH 7.8 - 8.3
–Baking soda: pH 9.0
Acids & Bases
•Acids
–Donate H+ when dissolved in water
–Acidic solutions have pH < 7
•Bases
–Accept H+ when dissolved in water
–Acidic solutions have pH > 7
Weak and Strong Acids
•Weak acids
–Reluctant H+ donors
–Can also accept H after giving it up
–Carbonic acid (H2CO3) is example
•Strong acids
–Completely give up H+ when dissolved
–Hydrochloric acid (HCl) is example
Buffer Systems
•Minimize shifts in pH
•Partnership between weak acid andbase it forms when dissolved
•Two work as pair to counter shifts in pH
Carbonic Acid-BicarbonateBuffer System
•When blood pH rises, carbonic aciddissociates to form bicarbonate and H+
H2C03 -----> HC03- + H+
•When blood pH drops, bicarbonate bindsH+ to form carbonic acid
HC03- + H+ -----> H2C03
Salts
•Compounds that release ions other thanH+ and OH- when dissolved in water
•Example: NaCl releases Na+ and Cl–
•Many salts dissolve into ions that playimportant biological roles