1. It’s Alive! Marieb Chapter 2: Part A (Chemistry Comes Alive)
Student Version
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2. Major concepts in this chapter What are they?
Energy vs. Matter
Major concepts in this chapter
What are they?
How do they work in the body?
Lets compare their properties:
Energy
Matter • - •
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3. Matter—anything that has mass and occupies space
3 states of matter
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4. Energy can be converted from
Capacity to do work
Types of energy
—energy in action
—stored (inactive) energy
Energy can be converted from
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5. Forms of Energy energy Stored in bonds of chemical substances
Results from movement of charged particles
Directly involved in moving matter
Travels in waves (e.g., visible light, ultraviolet light, and x-rays)
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6. Energy may be converted from one form to another
Energy Conversions
Energy may be converted from one form to another
Energy conversion is inefficient
Some energy is “lost” as heat (partly unusable energy)
Gasoline is ___________ energy; we can only remove % of the available energy to do work
Lance Armstrong’s muscles are 23% efficient
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7. The Energy Name Game! A ? B ?
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8. Name the Energy!
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9. Potential vs Kinetic Energy … Ole!
Video
Potential vs Kinetic Energy … Ole!
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10. Challenge Yourself! What Type of Energy?
Potential or Kinetic?
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11. Composition of Matter: Elements
Matter is composed of elements
Elements cannot be broken into simpler substances by ordinary chemical methods
Each has its own unique properties
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12. Composition of Matter Atoms Atomic symbol
We will use the term atom rather than element
Atomic symbol
One- or two-letter chemical shorthand for each element
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13. The Periodic Table Lanthanides Actinides
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14. Major Elements of the Human Body
Four elements make up ?% of body mass
Element
Carbon
Hydrogen
Oxygen
Nitrogen
Atomic symbol C H O N
Know these symbols!
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15. Lesser Elements of the Human Body
9 elements make up 3.9% of body mass
Element
Calcium
Phosphorus
Potassium
Sulfur
Sodium
Chlorine
Magnesium
Iodine
Iron
Atomic symbol Ca P K S Na Cl Mg I Fe
Yep, you need to know these too!
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16. Trace Elements of the Human Body
Occur in very minute amounts
11 elements make up < 0.01% of body mass
Many are part of, or activate, enzymes
For example:
Element
Chromium
Copper
Fluorine
Manganese
Silicon
Zinc
Atomic symbol Cr Cu F Mn Si Zn
These are important but you won’t find them
mentioned as often. Don’t worry about these symbols.
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17. How Important Are These Elements?
This is a generic form of _____________
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18. © 2013 Pearson Education, Inc.
Table 2.1.1

19. © 2013 Pearson Education, Inc.
Table 2.1.2

20. Atoms are composed of subatomic particles
Atomic Structure
Atoms are composed of subatomic particles
Protons and neutrons are in the nucleus
Electrons orbit the nucleus in an electron cloud
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21. Atomic Structure: The Nucleus
Almost entire mass of the atom
Neutrons
Carry no charge (Neutral!)
Protons
Carry positive charge
Mass = 1 amu
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22. Atomic Structure: Electrons
Electrons in orbitals within electron cloud
Circle outside the nucleus
Carry negative charge
Much smaller than protons or neutrons
Number of protons and electrons are always equal
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23. Figure 2.1 Two models of the structure of an atom.
Nucleus
Nucleus
Helium atom
Helium atom
2 protons (p+)
2 protons (p+)
2 neutrons (n0)
2 neutrons (n0)
2 electrons (e–)
2 electrons (e–)
Planetary model
Orbital model
Proton
Neutron
Electron
Electron
cloud
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24. Different atoms contain different numbers of subatomic particles
Identifying Elements
Different atoms contain different numbers of subatomic particles
Hydrogen has 1 proton, 0 neutrons, and 1 electron
Lithium has 3 protons, 4 neutrons, and 3 electrons
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25. Figure 2.2 Atomic structure of the three smallest atoms.
Proton
Neutron
Electron
Hydrogen (H)
(1p+; 0n0; 1e–)
Helium (He)
(2p+; 2n0; 2e–)
Lithium (Li)
(3p+; 4n0; 3e–)
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26. Combining Matter: Forming Molecules
Most atoms chemically combined with other atoms to form molecules
Molecule
Two or more atoms bonded together (e.g., H2 or C6H12O6)
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27. Most matter exists as mixtures Three types of mixtures
Two or more components physically intermixed
Three types of mixtures
Solutions
Colloids
Suspensions
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28. Types of Mixtures: Solutions
Homogeneous mixtures
Most are true solutions in body
Gases, liquids, or solids dissolved in water
Usually transparent, e.g., atmospheric air or salt solution
Solvent
Usually a liquid; usually water
Solute(s)
For example:
If glucose is dissolved in blood, glucose is the ; blood is the
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29. How Do We Make Solutions?
Can be expressed as
Percent of solute in total solution
(solvent assumed to be water)
Parts solute per 100 parts solvent
Milligrams per deciliter (mg/dl)
(weight per volume)
Molarity, or moles per liter (M)
1 mole of an element or compound = Its atomic or molecular weight (sum of atomic weights) in grams
1 mole of any substance contains 6.02  1023 molecules of that substance (Avogadro’s number)
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30. Colloids and Suspensions
Colloids (also called emulsions)
Heterogeneous mixtures, e.g., cytosol
Large solute particles do not settle out
Suspensions
Heterogeneous mixtures
Large, visible solutes settle out
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31. Figure 2.4 The three basic types of mixtures
Solution
Colloid
Suspension
Solute particles are larger than
in a solution and scatter light;
do not settle out.
Solute particles are very large,
settle out, and may scatter light.
Solute particles are very tiny,
do not settle out or scatter light.
Solute
particles
Solute
particles
Solute
particles
Example
Example
Jello
Whole
Blood
Example
Mineral water
Plasma
Settled red
blood cells
Unsettled
Settled
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32. What About Those Electrons?
Electrons have potential energy
Each shell corresponds to a specific level of potential energy
Shells that are farther from the nucleus contain electrons with more potential energy
Up to seven electron shells (energy levels) can exist
Atoms are most stable when their outermost (valence) shell is full
Atoms will interact with other atoms to fill their outermost shells (via chemical bonds)
Octet rule (rule of eights)
Except for the first shell (full with two electrons) atoms interact to have eight electrons in their valence shell
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33. Chemically Inert Elements
Valence shell fully occupied or contains eight electrons
Stable and unreactive because of this
The “loners”
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34. Figure 2.5a Chemically inert and reactive elements.
Chemically inert elements
Outermost energy level (valence shell) complete 8e 2e 2e
Helium (He)
(2p+; 2n0; 2e–)
Neon (Ne)
(10p+; 10n0; 10e–)
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35. Chemically Reactive Elements
Outer shell not full
Tend to gain, lose, or share electrons (form bonds) with other atoms to achieve stability
The “party animals” - like to interact with other atoms
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36. Figure 2.5b Chemically inert and reactive elements.
Chemically reactive elements
Outermost energy level (valence shell) incomplete 4e 1e 2e
Hydrogen (H)
(1p+; 0n0; 1e–)
Carbon (C)
(6p+; 6n0; 6e–) 1e 6e 8e 2e 2e
Oxygen (O)
(8p+; 8n0; 8e–)
Sodium (Na)
(11p+; 12n0; 11e–)
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37. Types of Chemical Bonds
We will discuss three major types:
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38. Attraction of opposite charges results in an ionic bond
Ionic Bonds
Ions are formed
An atom gains or loses electrons and becomes charged
The transfer of valence shell electrons from one atom to another forms ions
One becomes an anion (negative charge)
Atom that gained one or more electrons
One becomes a cation (positive charge)
Atom that lost one or more electrons
Attraction of opposite charges results in an ionic bond
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39. Figure 2.6a–b Formation of an ionic bond.
(17p+; 18n0; 18e–))
(11p+; 12n0; 10e–)) — +
Sodium atom (Na)
(11p+; 12n0; 11e–)
Chlorine atom (Cl)
(17p+; 18n0; 17e–)
Sodium ion (Na+)
Chloride ion (Cl–)
Sodium chloride (NaCl)
Sodium gains stability by losing
one electron, and chlorine becomes
stable by gaining one electron.
After electron transfer,
the oppositely charged ions
formed attract each other.
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40. Most ionic compounds are salts
When dry, salts form crystals instead of individual molecules
Example is NaCl (sodium chloride)
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41. Figure 2.6c Formation of an ionic bond.
Cl–
Na+
Large numbers of Na+ and Cl– ions
associate to form salt (NaCl) crystals.
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42. Do Salts Exist In the Body?
Rarely, because most dissolve in water!
Most ionic bonds are easily broken by water; ions are released
Their ions do exist in the body =
Exceptions:
Calcium phosphate in bone
Kidney stones (ouch!)
Gall stones (also ouch!)
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43. most salt crystals dissolve in water
Ionic Bonds are Weak -
most salt crystals dissolve in water
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Figure 2.12

44. Common Ions You Need To Know
OH- Hydroxide
PO43- Phosphate
HCO3- Bicarbonate
NH4+ Ammonium
H+ Hydrogen (proton)
Na+ Sodium
K+ Potassium
Ca2+ Calcium
Cl- Chloride
I- Iodide
COO- Carboxyl
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45. Formed by of two or more valence shell electrons
Covalent Bonds
Formed by of two or more valence shell electrons
Allows each atom to fill its valence shell at least part of the time
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46. Figure 2.7a Formation of covalent bonds.
Reacting atoms
Resulting molecules + or
Structural formula
shows single bonds.
Hydrogen atoms
Carbon atom
Molecule of methane gas (CH4)
Formation of four single covalent bonds:
Carbon shares four electron pairs with
four hydrogen atoms.
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47. Figure 2.7b Formation of covalent bonds.
Reacting atoms
Resulting molecules + or
Structural formula
shows double bond.
Oxygen atom
Oxygen atom
Molecule of oxygen gas (O2)
Formation of a double covalent bond: Two
oxygen atoms share two electron pairs.
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48. Figure 2.7c Formation of covalent bonds.
Reacting atoms
Resulting molecules + or
Structural formula
shows triple bond.
Nitrogen atom
Nitrogen atom
Molecule of nitrogen gas (N2)
Formation of a triple covalent bond: Two
nitrogen atoms share three electron pairs.
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49. Two Different Types of Covalent Bonds
Non-polar Covalent bonds
Electrons equally shared
Polar Covalent Bonds
Electrons UNequally shared
Come in single, double, or triple forms
More lines means more electrons shared
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50. Single, Double, and Triple Covalent Bonds
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51. Chemical and Structural Formulas
C2H5OH
Chemical Formula
(chemist’s shorthand)
Structural Formula
Sometimes you only see the “skeleton” of a molecule!
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52. What Do I Mean?
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53. More “Shorthand” - The Missing Atoms
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54. More “Shorthand” - The Missing Atoms
This is Cholesterol
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55. Here Are The Missing Atoms!
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56. Nonpolar Covalent Bonds
Electrons shared equally
Which bonds are nonpolar?
C - C
C - H
If we have a lot of these bonds in a biological molecule, the molecule is
non-polar (Most lipids= fats!)
It would be oily, greasy, and hate water
(hydrophobic = )
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57. Polar Covalent Bonds Unequal sharing of electrons produces polar bonds
Examples of polar covalent bonds:
C - O
N - H
O - O
N - O
O - H
A molecule with lots of polar covalent bonds loves to interact with the polar covalent bonds in water
Hydrophilic =
Most biological molecules are polar.
Carbohydrates • Some lipids
Proteins
Nucleic acids
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58. Water Is A Polar Moleculeδ– δ– δ+ δ+ δ+ δ+
V-shaped water (H2O) molecules have two
poles of charge—a slightly more negative
oxygen end (δ–) and a slightly more positive
hydrogen end (δ+).
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59. Figure 2.9 Ionic, polar covalent, and nonpolar covalent bonds compared.
Ionic bond
Polar covalent
bond
Nonpolar
covalent bond
Complete
transfer of
electrons
Unequal sharing
of electrons
Equal sharing of
electrons
Separate ions
(charged
particies)
form
Slight negative
charge (δ–) at
one end of
molecule, slight
positive charge (δ+)
at other end
Charge balanced
among atoms δ– δ+ δ+
Sodium chloride
Water
Carbon dioxide
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60. Let’s Watch A Video or Two or Three…!
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61. Hydrogen Bonds
Attractive force between a + charged hydrogen atom of one molecule and an charged atom of another molecule
Common between molecules with polar bonds such as water
Also found within a molecule, holding it in a three-dimensional shape
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62. Figure 2.10a Hydrogen bonding between polar water molecules.δ+ δ+ δ−
Hydrogen bond
(indicated by
dotted line) δ+ δ+ δ+ δ− δ− δ+ δ− δ+ δ+ δ+ δ+ δ−
The slightly positive ends (δ) of the water molecules
become aligned with the slightly negative ends (δ−)
of other water molecules.
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63. Figure 2.10b Hydrogen bonding between polar water molecules causes water to be “strong”.
A water strider can walk on a pond because of the high
surface tension of water, a result of the combined
strength of its hydrogen bonds.
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64. DNA Has Lots Of Hydrogen Bonds
Sugar:
Deoxyribose
Base:
Adenine (A)
Phosphate
Thymine (T)
Sugar
Phosphate
Adenine nucleotide
Thymine nucleotide
Hydrogen
bond
Deoxyribose
sugar
Sugar-
phosphate
backbone
Phosphate
Adenine (A)
Thymine (T)
Cytosine (C)
Guanine (G)
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65. Water vs. Ice
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66. Hydrogen Bond Videos http://www.youtube.com/watch?v=LK7ERiCy5b8
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67. Chemical Groups In Biological Molecules
Hydroxyl
Amino
Phenyl
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68. Chemical Groups In Biological Molecules
Methyl
Carboxyl
Phosphate
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69. Chemical Groups In Biological Molecules
Hydrocarbon chain
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70. Polar Or Non-polar Molecule?
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71. Polar Or Non-polar Molecule?
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72. Polar Or Non-polar Molecule?
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73. Occur when chemical bonds are formed, rearranged, or broken
Chemical Reactions
Occur when chemical bonds are formed, rearranged, or broken
Represented as chemical equations using molecular formulas
Chemical equations contain
Reactant(s) = what you start with
Chemical composition of the product(s) = what you end up with
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74. Examples of Chemical Equations
Reactant(s)
H + H 
4H + C 
Product(s)
H2 (Hydrogen gas)
CH4 (Methane)
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75. Types of Chemical Reactions
Synthesis (combination) reactions
Decomposition reactions
Exchange (swapping) reactions
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76. Synthesis Reactions A + B  A-B Example:
Atoms or molecules combine to form larger, more complex molecule
Always involve bond formation
Anabolic reaction
SMALLER TO BIGGER
Example:
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77. Figure 2.11a Patterns of chemical reactions.
Synthesis reactions
Smaller particles are bonded
together to form larger,
more complex molecules.
Example
Amino acids are joined together to
form a protein molecule.
Amino acid
molecules
Peptide bond
Protein
molecule
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78. Decomposition Reactions
A - B  A + B
Molecule is broken down into smaller molecules or its constituent atoms
Reverse of synthesis reactions
Involve breaking of bonds
Catabolic reaction
BIGGER TO SMALLER
Example:
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79. Figure 2.11b Patterns of chemical reactions.
Decomposition reactions
Bonds are broken in larger
molecules, resulting in smaller,
less complex molecules.
Example
Glycogen is broken down to release
glucose units.
Glycogen
Glucose
molecules
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80. Exchange Reactions A-B + C  A-C + B
Also called displacement reactions
Involve both synthesis and decomposition
Bonds are both made and broken
SWAPPING ATOMS
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81. Figure 2.11c Patterns of chemical reactions.
Exchange reactions
Bonds are both made and broken
(also called displacement reactions).
Example
ATP transfers its terminal phosphate
group to glucose to form glucose-
phosphate. +
Adenosine triphosphate (ATP)
Glucose +
Adenosine diphosphate
(ADP)
Glucose-
phosphate
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82. Oxidation-Reduction (Redox) Reactions
Let’s talk about these later when we discuss metabolism in detail…
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83. Reversibility of Chemical Reactions
Almost all biochemical reactions are reversible
A + B  AB
AB  A + B
Chemical equilibrium occurs when a reaction is reversible
Many biological reactions appear irreversible
Due to energy requirements
Due to removal of products
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84. Rate of Chemical Reactions
Affected by
 Temperature   Rate
 Concentration of reactant   Rate
 Particle size   Rate
Catalysts:  Rate without being chemically changed or part of product
Enzymes are biological catalysts
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85. The Law of Mass Action
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86. The Law of Mass Action
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87. The Law of Mass Action
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