Chapter 2 – Chemistry
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Transcript of Chapter 2 – Chemistry
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Chapter 2 – Chemistry
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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
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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
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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”
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The Periodic Table of Elements
Table that lists all known elements and their important information
The elements are organized in a specific way
• Table
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• 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!!!
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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 __________
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Atoms
– simplest particles of an element that retain all the properties of that element.
Properties of atoms determine the
properties of matter they compose
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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
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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
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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
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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
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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!!!
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Filling Energy Levels
• Filling rules – 2,8,8 electrons
• Orbitals - probability
Yellow = nucleusBlue = level 1 = 2Red = level 2 = 8Green = level 3 = 8
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1
2
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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
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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
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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
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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
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Types Of Chemical Bonds, Overview
• https://www.youtube.com/watch?v=_M9khs87xQ8
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• 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
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Valence electrons = those in the outermost energy level
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1. Ionic Bonds
Bonding produced when atoms transfer electrons.
Not as strong as covalent bonds.
Produce charged atoms (ions).
•i.e. NaCl
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Na+ Cl-
----NaCl
•Na has lost an electron (thus the + charge)•Cl has gained an electron (thus the – charge)
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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
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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)
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Molecule
– simplest part of a substance that retains all of the properties of the substance
– only covalent compounds form
molecules
water
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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
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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).
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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)
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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)
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“THE WINTER BALL”
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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
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• Endergonic – net absorption of free energy– Temperature decreases to indicate an
absorption of energy
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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
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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
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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
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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
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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)
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• 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.
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Concentration
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What Makes Water Such a Good Solvent?
• The chemical nature of water is called POLARITY
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The Hydrogen Bondhttp://www.youtube.com/watch?v=aH2IbYs_XjY
Occurs between H and O and between H and N
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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
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Acids and Bases
• Acid – sour, corrosive - lemons
• Alkaline – bitter, smooth - bleach
• Chemical significance???
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• 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
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Dissociation of Salt in Waterhttp://www.youtube.com/watch?v=CLHP4r0E7hg
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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)
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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
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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)
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pH
• – Scale from 0 – 14 to show how acidic or alkaline a solution is
• Logarithmic scale (10 fold >/< for each step)
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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
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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