Chapter 2 pt 1hhh.gavilan.edu/jcrocker/documents/Ch02_lecturept1.pdf · • example NaCl dropped...

Copyright © 2009 Pearson Education, Inc..

Including the lecture Materials of

Gregory AhearnUniversity of North Florida

with amendments andadditions by

John Crocker

Chapter 2 pt 1

Atoms, Molecules, and Life

Copyright © 2009 Pearson Education Inc.

Question

Research

Hypothesis

M & M/Data

Results:Experiment or

additional observation

Results does not support hypothesis; revise hypothesis or

pose new one

Results supports hypothesis; make

additional predictions and test them

Primary steps of the scientific method

• Feedback• falsifiable


If a hypothesis is correct, when we test it, we can expect a particular outcome

We try to disprove hypothesis.

Control groups are tested along with experimental groups to provide a comparison of results


Snowberry fly mimicking a jumping spider

Figure 1.3Cx


Case study: spider mimicry

Poun

ce ra

te (%

of t

rials

in

whi

ch s

pide

r jum

ped

on fl

y)

Control group (untreated flies)

Experimental group (wing markings masked)

Figure 1.3D


BiosphereEcosystem

-abiotic factors-all organisms

CommunityPopulationOrganism

ECOSYSTEM LEVEL Eucalyptus forest

COMMUNITY LEVEL All organisms in eucalyptus forest

POPULATION LEVEL Group of flying foxes

ORGANISM LEVEL Flying fox

ORGAN SYSTEM LEVEL Nervous system

ORGAN LEVEL Brain

Brain Spinal cord

Nerve

TISSUE LEVEL Nervous

tissue

CELLULAR LEVEL Nerve cell

MOLECULAR LEVEL Molecule of DNA Figure 1.1


Organisms are made up of:organ systemsorganstissuescellsmolecules

ECOSYSTEM LEVEL Eucalyptus forest

COMMUNITY LEVEL All organisms in eucalyptus forest

POPULATION LEVEL Group of flying foxes

ORGANISM LEVEL Flying fox

ORGAN SYSTEM LEVEL Nervous system

ORGAN LEVEL Brain

Brain Spinal cord

Nerve

TISSUE LEVEL Nervous

tissue

CELLULAR LEVEL Nerve cell

MOLECULAR LEVEL Molecule of DNA Figure 1.1


Each level of organization builds on the one below it

At each level, new properties emerge

ATOMS AND MOLECULES

Biological function starts at the chemical level


2.1 What Are Atoms?

Elements:

substances that cannot be broken down by ordinary chemical means (ex/ carbon)

all atoms belong to one of 96 types of naturally occurring elements

life requires about 25 of these elements


2.1 What Are Atoms?

Atoms:

basic structural unit of matter

consist of charged particles

protons (+)

neutrons (0)

electrons (-)

smallest particle of an element

each element has a unique number of protons (atomic number)


Isotopes - Atoms of the same element with different numbers of neutrons

Radioactive isotopes –

spontaneously break apart

forming different kinds of atoms

releasing energy in the process.

Example: radioactive uranium isotopes decay and form lead in the process


Atoms are electrically neutral because they have and equal number of positive protons and negative electrons

Helium atom

2

2

2

Protons

Neutrons

Electrons

Nucleus


HYDROGEN (H)Atomic number = 1

CARBON (C)Atomic number = 6

NITROGEN (N)Atomic number = 7

OXYGEN (O)Atomic number = 8

Electron

Outermost electron shell (can hold 8 electrons)

First electron shell (can hold 2 electrons)

Electrons are arranged in shells

Electrons orbit around atomic nuclei at specific distances called electron shells

the outermost shell determines the chemical properties of an atom


Electrons can move from electron shell to electron shell.

Electrons move from an inner to an outer shell when absorbing energy.

Electrons move from an outer shell to an inner shell when releasing energy.

All life depends on this energy.

The energy boosts the electron to a higher-energy shell

The electron drops back into lower-energy shell, releasing energy as light

energy

light

12

3

An electron absorbs energy

–

+ +

–

+

–


Energy Capture and Release

Life depends on electrons capturing and releasing energy

Electron shells correspond to energy levels

Energy exciting an atom causes an electron jump from a lower- to higher-energy shell

Later, the electron falls back into its original shell, releasing the energy


2.2 How Do Atoms Form Molecules?

Molecules: two or more atoms of one or more elements held together by interactions among their outermost electron shells

Atoms interact with one another according to two basic principles:

An inert atom will not react with other atoms when its outermost electron shell is completely full or empty.

A reactive atom will react with other atoms when its outermost electron shell is only partially full.


Atoms Interact

Atoms will react with other atoms if the outermost shell is partially full (such atoms considered reactive)

Example: Oxygen, with 6 electrons in outermost shell (can hold 2 more electrons)


Atoms Interact

Reactive atoms gain stability by electron interactions (chemical reactions)

Electrons can be lost to empty the outermost shell

Electrons can be gained to fill the outermost shell

Electrons can be shared with another atom where both atoms have full outermost shells


Atoms Interact

Hydrogen and oxygen atoms gain stability by interacting with each other

Single electrons from each of two hydrogen molecules fill the outer shell of an oxygen atom


Atoms combine with each other to fill outer electron shells (e.g. hydrogen and oxygen have unfilled outer electron shells, and thus, can combine to form the water molecule).

The water molecule, with a filled outer electron shell, is more stable than either the hydrogen or oxygen atoms that gave rise to it.

The results of losing, gaining, or sharing electrons are chemical bonds—attractive forces that hold atoms together in molecules.



A molecule may be depicted in several ways.

Fig. 2-4

(a) All bonds shown

(b) Bonds within common groups omitted

(c) Carbons and their attached hydrogens omitted

(d) Overall shape depicted

CH3 CH2 CH2 CH2 OH

OH

CH

H

H

C

H

H

C

H

H

C O H

H

H


Types of bonds

Ionic bonds: formed by passing an electron from one atom to another

One partner becomes positive, the other negative, and they attract one another.

Na+ + Cl– becomes NaCl (sodium chloride)

Positively or negatively charged atoms are called ions.

+ cation

- anion


Ions and Ionic Bonds

Atoms that have lost electrons become positively charged ions (e.g. sodium: Na+)

Atoms that have gained electrons become negatively charged ions (e.g. chlorine: Cl-)

Oppositely charged ions are attracted to each other are bound into a molecule by ionic bonds


Ions and Ionic Bonds

Salt crystals are repeated, orderly arranged sodium and chloride ions


Types of bonds (continued)

Covalent bonds: bond between two atoms that share electrons in their outer electron shell

For example, an H atom can become stable by sharing its electron with another H atom, forming H2 gas.


Covalent Bonds

Atoms with partially full outer electron shells can share electrons

Two electrons (one from each atom) are shared in a covalent bond


Covalent Bonds

Covalent bonds are found in H2 (single bond), O2 (double bond), N2 (triple bond) and H2 O

Covalent bonds are stronger than ionic bonds but vary in their stability

Most biological molecules contain covalent bonds


Covalent bonds produce either nonpolar or polar molecules.

Nonpolar molecule: atoms in a molecule equally share electrons that spend equal time around each atom, producing a nonpolar covalent bond


Nonpolar covalent bonding in hydrogen

(uncharged)

Electrons spend equal time near each nucleus

Same charge on both nuclei

(a)

++––

Nonpolar covalent bonding in hydrogen

Fig. 2-6a


Polar Covalent Bonds

Atoms within a molecule may have different nuclear charges

Those atoms with greater positive nuclear charge pull more strongly on electrons in a covalent bond

In diatomic molecules like H2 , both atoms exert the same pulling force on bond electrons: the covalent bond is nonpolar



In molecules where atoms of different elements are involved (H2 O), the electrons are not always equally shared: these covalent bonds are polar



A molecule with polar bonds may be polar overall

H2 O is a polar molecule

The (slightly) positively charged pole is around each hydrogen

The (slightly) negatively charged pole is around the oxygen


Hydrogen bonds: weak electrical attraction between positive and negative parts of polar molecules

Example: the negative charge of oxygen atoms in water molecules attract the positive charge of hydrogen atoms in other water molecules

Hydrogen Bonds


Hydrogen Bonds

Polar molecules like water have partially charged atoms at their ends

Hydrogen bonds form when partial opposite charges in different molecules attract each other

The partially positive hydrogens of one water molecule are attracted to the partially negative oxygen on another


Hydrogen Bonds

Polar biological molecules can form hydrogen bonds with water, each other, or even within the same molecule

Hydrogen bonds are rather weak but can collectively be quite strong


Hydrogen bonds

Fig. 2-7

hydrogen bonds

O (–)

H (+)

H (+)

O (–)

H (+)

H (+)


Free Radicals

Some cellular reactions produce free radicals

Free radical: a molecule whose atoms have one or more unpaired electrons in their outer shells

Free radicals are highly unstable and reactive

Free radicals steal electrons, destroying other molecules

Cell death can occur from free radical attack


Free Radicals

Free radicals are involved in causing heart disease, Alzheimer’s, cancer, and aging

Antioxidants like vitamins C and E can render free radicals harmless


2.3 Why Is Water So Important To Life?

Water interacts with many other molecules.• Oxygen released by plants during

photosynthesis comes from water.• Water is used by animals to digest food. • Water is produced in chemical reactions that

produce proteins, fats, and sugars.


Many molecules dissolve easily in water.• Water is an excellent solvent, capable of

dissolving a wide range of substances because of its positive and negative poles.

• example NaCl dropped into H2 O• The positive end of H2 O is attracted to Cl–. • The negative end of H2 O is attracted to Na+. • These attractions tend to pull apart the

components of the original salt.


Water as a solvent

Fig. 2-8

Cl–

OCl–

Cl–

H

H

Na+

Na+

Na+


• Water-insoluble molecules are hydrophobic• Water molecules repel and drive together

uncharged and nonpolar molecules like fats and oils

• The “clumping” of nonpolar molecules is called hydrophobic interaction


• Surface tension: water tends to resist being broken

• Cohesion: water molecules stick together

Fig. 2-9

Water Molecules Tend to Stick Together



Hydrogen bonding between water molecules produces high cohesion• Water cohesion explains how water

molecules can form a chain in delivering moisture to the top of a tree



Cohesion of water molecules along a surface produces surface tension• Fishing spiders and water striders rely on

surface tension to move across the surface of ponds



Water molecules stick to polar or charged surfaces in the property called adhesion• Adhesion helps water climb up the thin tubes

of plants to the leaves


Water can form ions.• Water dissociates to become H+ and OH–.

H2 O OH- + H+

• Acid solutions have more H+ (protons). • Alkaline solutions have more OH– (hydroxyl

ions).• A base is a substance that combines with H+,

reducing its concentration.• pH measures the relative amount of H+ and OH–

in a solution.

Acid, Basic, and Neutral Solutions


A water molecule is ionized.

Fig. 2-10

hydrogen ion (H+)

hydroxide ion (OH–)

water(H2 O)

+(+)(–)

O

HH

O H

H



Solutions where H+ > OH- are acidic• e.g. Hydrochloric acid ionizes in water:

HCl H+ + Cl-

• Lemon juice and vinegar are naturally produced acids



Solutions where OH- > H+ are basic• e.g. Sodium hydroxide ionizes in water:

NaOH Na+ + OH-

• Baking soda, chlorine bleach, and ammonia are basic



The degree of acidity of a solution is measured using the pH scale• pHs 0-6 are acidic (H+ > OH-)• pH 7 is neutral (H+ = OH-)• pH 8-14 is basic (OH- > H+)


Buffers Maintain Constant pH

• A buffer is a compound that accepts or releases H+ in response to pH change

• The bicarbonate buffer found in our bloodstream prevents pH change



• If the blood becomes too acidic, bicarbonate accepts (and absorbs) H+ to make carbonic acid

HCO3- + H+

H2 CO3

bicarbonate hydrogen ion carbonic acid



• If the blood becomes too basic, carbonic acid liberates hydrogen ions to combine with OH- to form water

H2 CO3 + OH- HCO3

- + H2 Ocarbonic acid hydroxide ion bicarbonate water


Water stabilizes temperature• Temperature reflects the speed of molecular

motion• Water has a high specific heat so it heats

up very slowly • It requires 1 calorie of energy to raise the

temperature of 1g of water 1oC• Because it heats up very slowly water

moderates the effect of temperature change• Very low or very high temperatures may

damage enzymes or slow down important chemical reactions


Water Stabilizes Temperature

Water requires a lot of energy to turn from liquid into a gas (heat of vaporization)• Evaporating water uses up heat from its

surroundings, cooling the nearby environment (as occurs during sweating)

• Because the human body is mostly water, a sunbather can absorb a lot of heat energy without sending her/his body temperature soaring


Water Stabilizes Temperature

Water requires a lot of energy to be withdrawn in order to freeze (heat of fusion)

Water freezes more slowly than other liquids


Water Forms an Unusual Solid: Ice

• Most substances become denser when they solidify from a liquid

• Water molecules spread apart slightly during the freezing process

• Because of this ice is less dense than liquid water


Water Forms an Unusual Solid: Ice• Ice floats in liquid water• Ponds and lakes freeze from the top

down and never freeze completely to the bottom


Frozen water floats (left) and frozen benzene sinks (right)

Figure 2.13x2


- Lower water is protected by the surface layer of ice.

–Life can survive in cold water underneath ice.

–Spring thaw pushes nutrient-rich bottom water to surface


Like no other common substance on earth, water naturally exists in all three physical states:

Figure 2.10B

• solid

• liquid

• gas


Figure 2.10Bx


Organic refers to molecules containing a carbon skeleton

Inorganic refers to carbon dioxide and all molecules without carbon

Organic vs. Inorganic in Chemistry

Chapter 2 pt 1hhh.gavilan.edu/jcrocker/documents/Ch02_lecturept1.pdf · • example NaCl dropped...

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