lecture 3: chemistry of life

Lecture 3: Chemistry of Life

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Chemical Benefits and Costs

• Understanding of chemistry provides fertilizers, medicines, etc.

• Chemical pollutants damage ecosystems

Bioremediation

Use of living organisms to withdraw harmful substances

from the environment

ElementsElements• Fundamental forms of matter

• Can’t be broken apart by normal means

• 92 occur naturally on Earth

Less than 12 occur on the examLess than 12 occur on the exam

Most Common Elements in Living Organisms

Oxygen

Hydrogen

Carbon

Nitrogen

What Are Atoms?

• Smallest particles that retain properties

of an element

• Made up of subatomic particles:

– Protons (+)

– Electrons (-)

– Neutrons (no charge)

Fig. 2.3, p. 22

HYDROGEN HELIUM

electron

proton

neutron

Hydrogen and Helium Atoms Hydrogen and Helium Atoms

Atomic NumberAtomic Number

• Number of protons

• All atoms of an element have the same atomic number

• Atomic number of hydrogen = 1

• Atomic number of carbon = 6

Mass NumberMass Number

Number of protons

+Number of neutrons

Isotopes vary in mass number

Atomic MassAtomic Mass

IsotopesIsotopes• Atoms of an element

with different numbers of neutrons (different mass numbers)

• Carbon 12 has 6 protons, 6 neutrons

• Carbon 14 has 6 protons, 8 neutrons

RadioisotopesRadioisotopes• Have an unstable

nucleus that emits energy and particles

• Radioactive decay transforms radioisotope into a different element

• Decay occurs at a fixed rate

Radioisotopes as TracersRadioisotopes as Tracers• Example: Tracer Drug Study

– How long does a drug stay in the patient?– Determine dose guidelines

• Compound synthesized with a radioisotope

• Emissions from the tracer can be detected with special devices– Track levels in the blood, urine and feces

• Following movement of tracers is useful in many areas of biology

High SensitivityHigh Sensitivity

Very Low DoseVery Low Dose

Other Uses of RadioisotopesOther Uses of Radioisotopes• Drive artificial pacemakers• Internal structure

– Thyroid and bone scans

• Radiation therapyEmissions from some radioisotopes can destroy cells. Some radioisotopes are used to kill small cancers.

Thyroid ScanThyroid Scan

• Measures health of thyroid by detecting radioactive iodine taken up by thyroid gland

normal thyroid enlarged cancerous

What Determines Whether What Determines Whether Atoms Will Interact?Atoms Will Interact?

ElectronsElectrons

• Carry a negative charge

• Repel one another

• Are attracted to protons in the nucleus

• Move in orbitals - volumes of space that surround the nucleus

Z

X

When all p orbitals are full

y

Electron OrbitalsElectron Orbitals

• Orbitals can hold up to two electrons

• Atoms differ in the number of occupied orbitals

• Orbitals closest to nucleus are lower energy and are filled first

Shell ModelShell Model

• First shell

– Lowest energy

– Holds 1 orbital with up to 2

electrons

• Second shell

– 4 orbitals hold up to 8

electrons

CALCIUM20p+ , 20e-

Electron VacanciesElectron Vacancies

• Unfilled shells make atoms likely to react

• Hydrogen, carbon, oxygen, and nitrogen all have vacancies in their outer shells

CARBON6p+ , 6e-

NITROGEN7p+ , 7e-

HYDROGEN1p+ , 1e-

Chemical Bonds, Molecules, Chemical Bonds, Molecules,

& Compounds& Compounds• Bond is union between electron

structures of atoms

• Atoms bond to form molecules

• Molecules may contain atoms of only one element - O2

• Molecules of compounds contain more than one element - H2O

Chemical BondsChemical Bonds

Electrostatic

Covalent

Metallic

1. Ionic Bonding1. Ionic Bonding

• One atom loses electrons, becomes positively charged ion

• Another atom gains these electrons, becomes negatively charged ion

• Charge difference attracts the two ions to each other

Ion FormationIon Formation

• Atom has equal number of electrons and protons - no net charge

• Atom loses electron(s), becomes positively charged ion

• Atom gains electron(s), becomes negatively charged ion

Formation of NaClFormation 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

SODIUMATOM11 p+

11 e-

SODIUMION

11 p+

10 e-

electron transfer

CHLORINEATOM17 p+

17 e-

CHLORINEION

17 p+

18 e-

Fig. 2.10a, p. 26

Formation of NaClFormation of NaCl

2. Covalent Bonding2. Covalent Bonding

Atoms share a pair or pairs of electrons to fill outermost shell

•Single covalent bond

•Double covalent bond

•Triple covalent bond

Two Flavors of Two Flavors of Covalent BondsCovalent Bonds

Non-polarNon-polar Covalent Covalent• Atoms share electrons

equally• Nuclei of atoms have

same number of protons

• Example: Hydrogen gas (H-H)

PolarPolar Covalent Covalent• Number of protons in

nuclei of participating atoms is NOT equal

• Molecule held together by polar covalent bonds has no NET charge

• Electrons spend more time near nucleus with most protons– Example: Water – Electrons more attracted

to O nucleus than to H nuclei

Example+

O

H H

slight negative charge at this end

slight positive charge at this end

molecule hasno net charge( + and - balanceeach other)

KEEP YOUR EYE ON THE ELECTRONS

Hydrogen BondingHydrogen Bonding

A bond by Hydrogen between two atoms

• Important for O and N

• Lets two electronegative atoms interact– The H gives one a net + and the other one

that is still – is attracted to it.

• The H proton becomes “naked” because its electron gets pulled away.

Hydrogen bond figure

- -

- + -

Like Charge Atoms Repel Each Other

Opposite Charge Atoms Attract Each Other

KEEP YOUR EYE ON THE ELECTRONS

onelargemolecule

anotherlargemolecule

a largemoleculetwistedbackonitself Fig. 2.12, p. 27

Examples of Hydrogen BondsExamples of Hydrogen Bonds

Properties of WaterProperties of Water

•Polarity

•Temperature-Stabilizing

•Cohesive

•Solvent

Water Is a Polar Water Is a Polar Covalent MoleculeCovalent Molecule

• Molecule has no net charge

• Oxygen end has a slight negative charge

• Hydrogen end has a slight positive charge

O

H H

O

H

HO

H

H

+ _

++

+

_

+

+

Liquid WaterLiquid Water

Hydrophilic & HydrophobicHydrophilic & HydrophobicSubstancesSubstances

• Hydrophilic substances– Polar– Hydrogen bond with water – Example: sugar

• Hydrophobic substances– Nonpolar– Repelled by water– Example: oil

Temperature-Stabilizing Temperature-Stabilizing EffectsEffects

• Water absorbs a lot more heat than oil before its temperature rises.

• Why?

• Much of the added energy disrupts hydrogen bonding rather than increasing the movement of molecules

Evaporation of WaterEvaporation of Water

• Large energy input can cause individual molecules of water to break free into air

• As molecules break free, they carry away some energy (lower temperature)

• Evaporative water loss is used by mammals to lower body temperature

Why Ice FloatsWhy Ice Floats

• In ice, hydrogen bonds lock molecules in a lattice

• Water molecules in lattice are spaced farther apart then those in liquid water

• Ice is less dense than water

Water CohesionWater Cohesion• Hydrogen bonding holds

molecules in liquid water together

• Creates surface tension

• Allows water to move as continuous column upward through stems of plants

Water Is a Good SolventWater Is a Good Solvent

• Ions and polar molecules dissolve easily in water

• When solute dissolves, water molecules cluster around its ions or molecules and keep them separated

Fig. 2.16, p. 29

Na+

Cl–

– –

––

––

–

––

– –

+ ++

+

+

+

+

+

+

+

+

++ +

+

+

+

+

Spheres of HydrationSpheres of Hydration

WaterWater

• Solvent- polar– Keeps ions in solution– Doesn’t dissolve membranes

• Heat management– Loosing heat– Holding heat– Density Changes

If it wasn’t ugly enough already: If it wasn’t ugly enough already:

Hydrogen Ions: HHydrogen Ions: H++

• Unbound protons

• Have important biological effects

• Form when water ionizes

The pH ScaleThe pH Scale

• Measures H+ concentration of fluid• Change of 1 on scale means 10X

change in H+ concentration

Highest H+ Lowest H+

0---------------------7-------------------14Acidic Neutral Basic

Examples of pHExamples of pHPure water is neutral with pH of 7.0

Acidic

Basic

Acids & BasesAcids & 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

BuffersBuffersMinimize shifts in pH

• When blood pH rises, carbonic acid dissociates to form bicarbonate and H+

H2C03 -----> HC03- + H+

• When blood pH drops, bicarbonate binds H+ to form carbonic acid

HC03- + H+ -----> H2C03

AcidosisAcidosis

AlkalosisAlkalosis

Carbonic Acid-Bicarbonate Buffer SystemCarbonic Acid-Bicarbonate Buffer System

Lecture 2:Chemistry of Life

Part 2

Organic Compounds

Hydrogen and other elements covalently bonded to carbon

Carbohydrates

Lipids

Proteins

Nucleic Acids

Carbon’s Bonding Behavior

• Outer shell of carbon has 4 electrons; can hold 8

• Each carbon atom can form covalent bonds with up to four atoms

Bonding Arrangements

• Carbon atoms can form chains or rings

• Other atoms project from the carbon backbone

Functional Groups

• Atoms or clusters of atoms that are covalently bonded to carbon backbone

• Give organic compounds their different properties

Examples of Functional Groups

Hydroxyl group - OH

Amino group - NH3+

Carboxyl group - COOH

Phosphate group - PO3-

Sulfhydryl group - SH

Types of Reactions

Functional group transfer

Electron transfer

Rearrangement

Condensation

Cleavage

Condensation Reactions

• Form polymers from subunits

• Enzymes remove -OH from one molecule, H from another, form bond between two molecules

• Discarded atoms can join to form water

Fig. 3.4a, p. 37

enzyme action at functional groups

CONDENSATION

Hydrolysis

• A type of cleavage reaction

• Breaks polymers into smaller units

• Enzymes split molecules into two or more parts

• An -OH group and an H atom derived from water are attached at exposed sites

enzyme action at functional groups

HYDROLYSIS

Fig. 3.4b, p. 37

Carbohydrates

Monosaccharides(simple sugars)

Oligosaccharides(short-chain carbohydrates)

Polysaccharides(complex carbohydrates)

Monosaccharides

• Simplest carbohydrates

• Most are sweet tasting, water soluble

• Most have 5- or 6-carbon backbone

Glucose (6 C) Fructose (6 C)

Ribose (5 C) Deoxyribose (5 C)

Two Monosaccharides

glucose fructose

Disaccharides

• Type of oligosaccharide

• Two monosaccharides covalently bonded

• Formed by condensation reaction

+ H2O

glucose fructose

sucrose

Polysaccharides

• Straight or branched chains of many sugar monomers

• Most common are composed entirely of glucose– Cellulose

– Starch (such as amylose)

– Glycogen

Cellulose & Starch

• Differ in bonding patterns between monomers

• Cellulose - tough, indigestible, structural material in plants

• Starch - easily digested, storage form in plants

Cellulose and Starch

Glycogen

• Sugar storage form in animals

• Large stores in muscle and liver cells

• When blood sugar decreases, liver cells degrade glycogen, release glucose

Chitin

• Polysaccharide

• Nitrogen-containing groups attached to glucose monomers

• Structural material for hard parts of invertebrates, cell walls of many fungi

• Most include fatty acids– Fats– Phospholipids– Waxes

• Sterols and their derivatives have no fatty acids

• Tend to be insoluble in water

Lipids

Fatty Acids

• Carboxyl group (-COOH) at one end

• Carbon backbone (up to 36 C atoms)

– Saturated - Single bonds between carbons

– Unsaturated - One or more double bonds

Three Fatty Acids

stearic acid oleic acid linolenic acid

Fats

• Fatty acid(s)

attached to

glycerol

• Triglycerides

are most

common

Phospholipids

• Main components of cell

membranes

Sterols and Derivatives

• No fatty acids

• Rigid backbone of four

fused-together carbon

rings

• Cholesterol - most

common type in

animals

Waxes

• Long-chain fatty acids linked to

long chain alcohols or carbon

rings

• Firm consistency, repel water

• Important in water-proofing

•Omega-6 fatty acids are the predominant polyunsaturated fatty acids (PUFAs) in the Western diet.

•The omega-6 and omega-3 fatty acids are metabolically distinct and have opposing physiologic functions.

•The increased omega-6/omega-3 ratio in Western diets most likely contributes to an increased incidence of heart disease and inflammatory disorders.

•Omega-3 PUFAs suppress cell mediated immune responses and reduce inflammation

Polyunsaturated Fatty Acids

Omega-3

Omega-6

•Bioactive Lipids•Made in all cells•Short range signaling•Eicosanoids?

•Prostaglandins•Inflammation and Pain Perception•Kidney Function•Bone Development•Reproductive Process

•Commercially Important•$4 BILLION/ Year spend on drugs to inhibit prostaglandin synthesis•Vioxx, Celebrex, Ibuprofen, Asprin

Lipids in Cell Signaling

PGE2

Amino Acid Structure

aminogroup

carboxylgroup

R group

Properties of Amino Acids

• Determined by the “R group”

• Amino acids may be:

– Non-polar

– Uncharged, polar

– Positively charged, polar

– Negatively charged, polar

Protein Synthesis

• Protein is a chain of amino acids linked

by peptide bonds

• Peptide bond

– Type of covalent bond

– Links amino group of one amino acid with

carboxyl group of next

– Forms through condensation reaction

Primary Structure

• Sequence of amino acids

• Unique for each protein

• Two linked amino acids = dipeptide

• Three or more = polypeptide

• Backbone of polypeptide has N atoms:

-N-C-C-N-C-C-N-C-C-N-

Protein Shapes

• Fibrous proteins

– Polypeptide chains arranged as strands or sheets

• Globular proteins

– Polypeptide chains folded into compact, rounded

shapes

• Primary structure influences shape in two main ways:– Allows hydrogen bonds to form between

different amino acids along length of chain

– Puts R groups in positions that allow them to interact

Primary Structure & Protein Shape

Secondary Structure

• Hydrogen bonds form between different parts of polypeptide chain

• These bonds give rise to coiled or extended pattern

• Helix or pleated sheet

Examples of Secondary Structure

Tertiary Structure

Folding as a

result

of interactions

between R

groups

heme group

coiled and twisted polypeptide chain of one globin molecule

Quaternary Structure

Some proteins

are made up of

more than one

polypeptide

chain

Hemoglobin

Polypeptides With Attached Organic Compounds

• Lipoproteins

– Proteins combined with cholesterol, triglycerides,

phospholipids

• Glycoproteins

– Proteins combined with oligosaccharides

Denaturation

• Disruption of three-dimensional shape

• Breakage of weak bonds

• Causes of denaturation:– pH

– Temperature

• Destroying protein shape disrupts function

A Permanent Wave

hair wrapped around cuticles

differentbridges form

bridgesbroken

• Sugar

– Ribose or deoxyribose

• At least one phosphate group

• Base

– Nitrogen-containing

– Single or double ring structure

Nucleotide Structure

Nucleotide Functions

• Energy carriers

• Coenzymes

• Chemical messengers

• Building blocks for

nucleic acids

ATP - A Nucleotide

three phosphate groups

sugar

base

• Composed of nucleotides

• Single- or double-stranded

• Sugar-phosphate backbone

Nucleic AcidsAdenineCytosine

DNA

• Double-stranded • Consists of four

types of nucleotides

• A bound to T• C bound to G

RNA

• Usually single strands

• Four types of nucleotides

• Unlike DNA, contains the base uracil in place of thymine

• Three types are key players in protein synthesis

• Normal metabolic products of one

species that can harm or kill a different

species

• Natural pesticides

– Compounds from tobacco

– Compounds from chrysanthemum

Natural Toxins

Synthetic Toxins

atrazine DDTmalathion

Negative Effects of Pesticides

• May be toxic to predators that help fight pests

• May be active for weeks to years

• Can be accidentally inhaled, ingested, or absorbed by humans

• Can cause rashes, headaches, allergic reactions

Producers Capture Carbon

Using photosynthesis, plants and other producers turn carbon dioxide and

water into carbon-based compounds

Atmospheric Carbon Dioxide

• Researchers have studied concentration of CO2 in air since the 1950s

• Concentration shifts with season– Declines in spring and summer when

producers take up CO2 for photosynthesis

CO2 and Global Warming

• Seasonal swings in CO2 increasing

• Spring decline starting earlier

• Temperatures in lower atmosphere increasing

• Warming may be promoting increased photosynthesis

Humans and Global Warming

• Fossil fuels are rich in carbon

• Use of fossil fuels releases CO2 into atmosphere

• Increased CO2 may contribute to global warming