Chapter 2 – Chemistry

Why Study Chemistry in Biology?

• Chemical changes in matter are the foundation for all life processes

• Living things are composed of the same kinds of matter that make up nonliving things

Chemical breakdown of human body is:• 65% Oxygen• 18% Carbon Makes up 96% of living

things• 10% Hydrogen• 3% Nitrogen• 1.5% Calcium• .35% Potassium• .25% Sulfur• .15% Sodium• .15% Chlorine• .05% Magnesium• .0004% Iron• .00004% Iodine• Traces of fluorine, silicon, manganese, zinc,

copper, aluminum, and arsenic

How are these elements all put together to make up the human body?

• Matter – – anything that occupies space and has …

• Mass – – quantity of matter an object has. (Weight is not the

same – downward force of gravity factored in)

• Elements – – pure substances that cannot be broken down

chemically into simpler kinds of matter. – 118 elements as of 2006. – N, C, H, O, P, S are important elements in Biology

2.1 “Composition of Matter”

The Periodic Table of Elements

Table that lists all known elements and their important information

The elements are organized in a specific way

• Table

• Atomic number = – the number of protons in

atom (= electrons also)

• Element symbol– (Carbon)

• Mass number= – # of protons + # of

neutrons– Atomic Mass = relative

average mass of element – a decimal

6

C12

12.01

ALL ELEMENTS ARE ELECTRICALLY NEUTRAL TO START!!!

Element symbols to be familiar with…1. Carbon __________2. Hydrogen __________3. Oxygen _________4. Nitrogen __________5. Calcium__________6. Phosphorus __________7. Potassium __________8. Sulfur __________9. Sodium __________10. Chlorine __________11. Magnesium _________12. Iron __________13. Iodine __________14. Fluorine__________15. Silicon __________16. Zinc __________17. Copper __________18. Aluminum __________19. Arsenic __________20. Manganese __________

Atoms

– simplest particles of an element that retain all the properties of that element.

Properties of atoms determine the

properties of matter they compose

Parts of an Atom1. Protons:

Positive electrical charge

Mass: 1AMU

Location: in nucleus

2. Neutrons:

No electrical charge

Mass: 1AMU

Location: in nucleus

3. Electrons:

Negative electrical charge

Mass: 1/2000 (so it is counted as 0 AMU)

Location: surrounding nucleus

  Element SymbolAtomic Number

Mass Number

Number of Protons

Number of

Neutrons

Number of

Electrons

# Electrons in ourter most

orbital

1 Carbon   6 12        

2   O         8  

3 Magnesium       12 12    

4 Sodium   11 23        

5   N            

6 Chlorine   17 35        

7   H         1  

8   S 16          

9 Iodine   53 127     53  

10   Si     14      

11   Ca     20      

12   P 15 31        

13 Iron   26     30    

14 Aluminum           13  

15   Mn     25      

16 Copper   29 64        

17 Zinc         35 30  

18 Arsenic As            

19   K     19      

20 Fluorine     19     9  

Isotopes

Element that have the same number of protons but different number of neutrons

They vary by their atomic mass and mass number

For example: C-12, C-13, C-14

The decimal you see on the PT = average of the relative amounts in nature of the various isotopes

Models of the atom

Bohr model – electrons appear to “orbit” the nucleus – aka Planetary Model

Electron cloud model - protons and neutrons concentrated in the nucleus and electrons occupying various energy levels around the nucleus – not sure where the electrons are at any time

How are the electrons arranged in the energy levels?

• First energy level – will get a maximum of 2 electrons

• Second energy level = 8 electrons• Third energy level = 8 electrons

Electrons “sit” in these levels in ONLY this order!!!

Filling Energy Levels

• Filling rules – 2,8,8 electrons

• Orbitals - probability

Yellow = nucleusBlue = level 1 = 2Red = level 2 = 8Green = level 3 = 8

1

2

Diagram the following atoms1. 1H1

2. 12C6 atomic number

3. 14N7

4. 16O8 mass number

5. 23Na11

6. 35Cl17

7. 39K19

8. 40Ca20

NOTE – THESE ARE ISOTOPE DESIGNATIONS

Why Do Atoms Combine?

• Atoms will combine chemically to produce compounds

• Compounds form due to arrangement of electrons in outermost energy level=

VALENCE ELECTRONS – Atoms are most stable when outer

energy level is filled• Chemical bonds are broken, atoms are rearranged, and

new chemical bonds are formed. • Chemical Bonds = attractive forces holding atoms

together ALL OF THESE CHANGES INVOLVE AN EXCHANGE OF ENERGY

When atoms combine you get…

Compounds are:

Pure substances made up of atoms of 2 or more different elements– i.e. Water, Glucose, Salt– can be ionic or covalent

Molecules are:

Pure substances made up of atoms of 2 or more similar elements, i.e. O2

- Can only be covalent

Atoms in molecules and compounds are arranged in fixed proportions

• The chemical formula:

2H2O

coefficient

CH4 (Methane)

– 1 C 4 H

• C6H12O6 (Glucose)

– 6 C 12 H 6 O

• (NH4)2SO4 (Ammonium Sulfate)

– 2 N 8H 1S 4 O

subscript

Types Of Chemical Bonds, Overview

• https://www.youtube.com/watch?v=_M9khs87xQ8

• Physical and Chemical properties of compounds differ from the elements that make them up.– i.e. NaCl, H2Ohttp://www.bing.com/videos/search?

q=ionic+and+covalent+bonds&qs=n&form=QBVR&pq=ionic+and+covalent+bonds&sc=8-24&sp=-1&sk=#view=detail&mid=FC661AB5D4927AD1FDD7FC661AB5D4927AD1FDD7

Valence electrons = those in the outermost energy level

1. Ionic Bonds

Bonding produced when atoms transfer electrons.

Not as strong as covalent bonds.

Produce charged atoms (ions).

•i.e. NaCl

Na+ Cl-

----NaCl

•Na has lost an electron (thus the + charge)•Cl has gained an electron (thus the – charge)

2. Covalent Bonds

Covalent Bonds – form when 2 atoms share 1 or more

electrons

H2O, CH4, C6H12O6

– Strong bonds– All organic (have C & H) substances

H2O

Covalent Bonds

Covalent bonds are strong bonds due to the sharing of electrons.

– In order to break a covalent bond, • High heat • Electrical current • High Pressure• Enzymes (catalysts produced by living things)

Molecule

– simplest part of a substance that retains all of the properties of the substance

– only covalent compounds form

molecules

water

Using the model kits, make the following*:

1. H2O2. O2

3. CH4

4. HCl5. O3

6. CO2

7. H2

8. C2H6

9. C2H4

10. C2H2

11. C3H8

12. NH3

* Draw Structural Diagrams of each* Draw Structural Diagrams of each

•Colors of spheres on board•Use wooden sticks for single bonds; springs for double and triple bonds

2.2 “Energy”Energy – Ability to do work or cause change

All living things need a constant flow of energy – types???

1st Law of Energy: Energy is neither created nor destroyed but transferred

Free energy – energy in a system that is available to do work.

Body contains glucose to provide free energy and stored energy (glycogen and fat). To tap into these, you must break these down (digestion) into their simplest form (glucose).

States of Matter

Solids, liquids and gases

Particle movement (greater in gases) – higher energy

Shape and volume (fixed in solids)

Concentration of particles (tighter in a solid)

Chemical Reactions

CO2 + H2O H2CO3 (Carbonic acid)

Reactants Product

Yields

The above reaction occurs in one direction and is non-reversible.

The reaction below occurs in both directions and is reversible.

CO2 + H2O H2CO3 (Carbonic acid)

“THE WINTER BALL”

Energy Transfer Body continuously goes through series of

chemical reactions = METABOLISM

– Exergonic – net release of free energy– Temperature increases to indicate a release of

energy

• Endergonic – net absorption of free energy– Temperature decreases to indicate an

absorption of energy

Activation Energy Amount of energy necessary to begin a

reaction

Catalysts – Chemicals that lower the amount of needed Activation Energy

Enzymes – Organic catalysts (found only in living things) – help without being changed

Lactase breaks down the milk sugar Lactose

Redox Reactions

INVOLVE Transfer of electrons

Oxidation reactions – Reactants lose one or more electrons forming + ions

• Reduction reactions – Reactants gain one or more electrons. Form – ions

• Na + Cl Na+Cl-• Oxidized Reduced

• These reactions always occur together

2.3 “Water and Solutions”• Most mass of living things is water

(Universal solvent)

–About 65% of the total mass of our cells is water

–Chemical reactions occur in water

–Must understand chemistry of water

1. Add water to a beaker, about half full2. Into another beaker, add the same amount of rubbing alcohol3. Into a third beaker, do the same with cooking oil4. Sprinkle a small amount of salt into each and swirl5. Let stand for a moment

• Name the solute and solvent in each beaker.

• Solute Solvent Salt Water, Alcohol, cooking oil

1. In which beaker did the salt dissolve (go into solution)?

1. The water and somewhat in alcohol

2. Which solution is an aqueous solution (one in which the solvent is water)?

1. Only the water

DEMO

Solutions

A mixture of one or more substances uniformly distributed in another substance – physically combined

• Solute – substance being dissolved (sugar)

• Solvent – substance that does the dissolving (water)

• Physically but not chemically combined

• Solutions can vary in concentration of solute; 5% sugar solution has 5% sugar and 95% water– A solution is said to be saturated when the

solvent can no longer dissolve all the solute

• Aqueous solution – solvent is water

Solutions, cont.

Concentration

What Makes Water Such a Good Solvent?

• The chemical nature of water is called POLARITY

The Hydrogen Bondhttp://www.youtube.com/watch?v=aH2IbYs_XjY

Occurs between H and O and between H and N

Properties of Water(due to its polar nature)

• Cohesion – water sticking to itself – a barrel of monkeys- surface tension

• http://www.youtube.com/watch?v=ynk4vJa-VaQ

• Adhesion – water sticking to another polar substance – glass slide demo - capillarity

• Thermal regulation – high heat capacity, evaporative cooling

• Density of ice

Acids and Bases

• Acid – sour, corrosive - lemons

• Alkaline – bitter, smooth - bleach

• Chemical significance???

• Ionization or Dissociation– the production of ions when atoms or

molecules break apart

NaCl Na+ and Cl- ionic dissociate

H20 H+ + OH – covalent ionize

• H+ = Hydrogen ion

• OH- = Hydroxide ion

• H3O+ = Hydronium ion

Dissociation of Salt in Waterhttp://www.youtube.com/watch?v=CLHP4r0E7hg

Production of the Hydronium ion• Due to the high kinetic energy of the

molecules of water, there are numerous collisions. Some of these collisions are strong enough to dislodge protons (H+) from a water molecule or from an ionized acid molecule such as H+Cl-.

• Other water molecules will pick up these stray protons

• H2O H+ + OH-

• H20 + H+ H3O+ (Hydronium ion)

Acids

• Acid Solution – # of H30+ ions outnumbers the OH- ions in a

solution

• HCl = H+ + Cl-

• H+ + H2O H3O+ Hydronium ion

• Acids are sour and corrosive• Acid rain - pH of normal rain ~ 5.0 – 5.6

on pH scale• SO3 + H2O H2SO4

Bases (Alkaline)

• Base Solution – # of OH- ions outnumbers the H+ ions in a

solution

• NaOH --- Na+ + OH-

• Alkaline solutions are bitter

• Feel slippery (OH- ions react with oils of skin forming a soap)

pH

• – Scale from 0 – 14 to show how acidic or alkaline a solution is

• Logarithmic scale (10 fold >/< for each step)

Activity1. Using universal pH paper, determine the pH of the

following solutions2. Dip the tip of the pH paper into the solution, wait a

minute and compare it to the colored scale on the vial

3. List the pH’s of the following solutions:a. ammoniab. vinegarc. milkd. black coffeee. baking soda solutionf. colag. milk of magnesiah. lemon juicei. water

4. On a blank pH scale, place the solutions in the proper spot

Buffers• Buffers

– Naturally control the pH in living systems

• Neutralize small amounts of acids and bases• Maintain homeostasis• Enzymes in body require a particular pH!!!

– Stomach acids and urine – acidic

• Blood and intestinal fluids are alkaline• Neutralization reactions– Occurs when acids &

bases react w/each other• Results in the formation of a salt and water

Na+OH- + H+Cl- NaCl + H2O