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Chapter 02
Lecture and Animation Outline
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1
The Nature of Molecules andthe Properties of Water
Chapter 2
2
3
Nature of Atoms
• Matter has mass and occupies space
• All matter is composed of atoms
• Understanding the structure of atoms is critical to understanding the nature of biological molecules
4
Atomic Structure
• Atoms are composed of– Protons
• Positively charged particles• Located in the nucleus
– Neutrons• Neutral particles• Located in the nucleus
– Electrons • Negatively charged particles• Found in orbitals surrounding the nucleus
5
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neutron(no charge)
1 Proton1 Electron
8 Protons8 Neutrons8 Electrons
Hydrogen Oxygen
proton(positive charge)
electron(negative charge)
a.
b.
6
Atomic number
• Number of protons equals number of electrons– Atoms are electrically neutral
• Atomic number = number of protons– Every atom of a particular element has the
same number of protons
• Element– Any substance that cannot be broken down to
any other substance by ordinary chemical means
7
Atomic mass
• Mass or weight?– Mass – refers to amount of substance– Weight – refers to force gravity exerts on
substance
• Sum of protons and neutrons is the atom’s atomic mass
• Each proton and neutron has a mass of approximately 1 Dalton
Electrons
• Negatively charged particles located in orbitals
• Neutral atoms have same number of electrons and protons
• Ions are charged particles – unbalanced– Cation – more protons than electrons = net
positive charge– Anion – fewer protons than electrons = net
negative charge8
Isotopes
• Atoms of a single element that possess different numbers of neutrons
• Radioactive isotopes are unstable and emit radiation as the nucleus breaks up– Half-life – time it takes for one-half of the
atoms in a sample to decay
9
10
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Carbon-12 Carbon-13 Carbon-14
6 Protons 6 Neutrons 6 Electrons
6 Protons 7 Neutrons 6 Electrons
6 Protons 8 Neutrons 6 Electrons
11
Electron arrangement
• Key to the chemical behavior of an atom lies in the number and arrangement of its electrons in their orbitals
• Bohr model – electrons in discrete orbits
• Modern physics defines orbital as area around a nucleus where an electron is most likely to be found
• No orbital can contain more than two electrons
12
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One spherical orbital (1s)
Electron Shell Diagram Corresponding Electron Orbital
EnergyLevel K
x
z
y
Neon
Electron Shell Diagram Electron Orbitals
x
y z
One spherical orbital (2s)
Corresponding Electron OrbitalsElectron Shell Diagram
Three dumbbell-shaped orbitals (2kk)
Energylevel L
Energy levels
• Electrons have potential energy related to their position– Electrons farther from nucleus have more
energy
• Be careful not to confuse energy levels, which are drawn as rings to indicate an electron’s energy, with orbitals, which have a variety of three-dimensional shapes and indicate an electron’s most likely location 13
14
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NM
LK
Energyabsorbed
Nucleus
Energyreleased
Nucleus
KLMN
15
Redox
• During some chemical reactions, electrons can be transferred from one atom to another– Still retain the energy of their position in the atom– Oxidation = loss of an electron– Reduction = gain of an electron
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++
Oxidation Reduction
+
16
Elements
• Periodic table displays elements according to valence electrons
• Valence electrons – number of electrons in outermost energy level
• Inert (nonreactive) elements have all eight electrons
• Octet rule – atoms tend to establish completely full outer energy levels
17
Periodic Table of the Elements
1
H
3Li
4Be
19
K
12
Mg
93
Np94
Pu95
Am96
Cm97
Bk98
Cf99
Es100
Fm101
Md102
No103
Lr
37
Rb38
Sr39
Y42
Mo45
Rh46
Pd47
Ag48
Cd49
In50
Sn51
Sb52
Te53
I54
Xe
21
Sc40
Zr
22
Ti23
V24
Cr25
Mn27
Co28
Ni29
Cu30
Zn36
Kr
5
B6 C
8
O 2
He
55
Cs56
Ba72
Hf73
Ta74
W75
Re76
Os77
Ir78
Pt79
Au80
Hg81
Tl82
Pb83
Bi84
Po85
At86
Rn87
Fr88
Ra
57
La89
Ac104 105 106 107 108 109
58
Ce59
Pr60
Nd61
Pm62
Sm63
Eu64
Gd65
Tb66
Dy67
Ho68
Er69
Tm70
Yb71
Lu
90
Th91
Pa92
U(Actinide series)
11
Na20
Ca
41
Nb43
Tc44
Ru
26
Fe
13
Al
31
Ga32
Ge
14
Si
7
N15
P
33
As
16
S35
Br34
Se
9
F
18
Ar
10
Ne17
Cl
110
Rf Ob Sg Bh Hs Mt Ds
111 112 113 114 115 116 117
Rg Cn Uut Fl Uup Lv Uus Uuo
118
Key
1
Hatomic number
chemical symbol
Mg
Fe
S
P
Ca
Cl
Na
N
H
C
a. b.
O
(Lanthanide series)
Carbon (C)
Oxygen (O)
Hydrogen (H)
Nitrogen (N)
Sodium (Na)
Chlorine (Cl)
Calcium (Ca)
Phosphorus (P)
Potassium (K)K
Sulfur (S)
Iron (Fe)
Magnesium (Mg)
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• 90 naturally occurring elements• Only 12 elements are found in
living organisms in substantial amounts
• Four elements make up 96.3% of human body weight– Carbon, hydrogen, oxygen,
nitrogen
• Organic molecules contain primarily CHON
• Some trace elements are very important
18
Mg
Fe
S
P
Ca
Cl
Na
N
H
C
b.
O
Carbon (C)
Oxygen (O)
Hydrogen (H)
Nitrogen (N)
Sodium (Na)
Chlorine (Cl)
Calcium (Ca)
Phosphorus (P)
Potassium (K)K
Sulfur (S)
Iron (Fe)
Magnesium (Mg)
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19
Chemical Bonds
• Molecules are groups of atoms held together in a stable association
• Compounds are molecules containing more than one type of element
• Atoms are held together in molecules or compounds by chemical bonds
20
Ionic bonds
• Formed by the attraction of oppositely charged ions
• Gain or loss of electrons forms ions– Na atom loses an electron to become Na+
– Cl atom gains an electron to become Cl–
– Opposite charges attract so that Na+ and Cl– remain associated as an ionic compound
• Electrical attraction of water molecules can disrupt forces holding ions together
21
Sodium atom
Cl
Chlorine atom
a.
Sodium ion (+)
Cl–
Chloride ion (–)
b. NaCl crystal
Cl– Cl–
Cl– Cl–
Cl–
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NaNa++
NaNa++
NaNa++
NaNa++
NaNa Na+Na+
22
Covalent bonds
• Form when atoms share 2 or more valence electrons
• Results in no net charge, satisfies octet rule, no unpaired electrons
• Strength of covalent bond depends on the number of shared electrons
• Many biological compounds are composed of more than 2 atoms – may share electrons with 2 or more atoms
23
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NN
N2N N
H H
O2O O
H2
O O
Single covalent bondHydrogen gas
covalent bond
Double covalent bondoxygen gas
Triple covalent bondNitrogen gas
H H
24
Electronegativity
• Atom’s affinity for electrons
• Differences in electronegativity dictate how electrons are distributed in covalent bonds– Nonpolar covalent bonds = equal sharing of
electrons– Polar covalent bonds = unequal sharing of
electrons
25
Chemical reactions
• Chemical reactions involve the formation or breaking of chemical bonds
• Atoms shift from one molecule to another without any change in number or identity of atoms
• Reactants = original molecules• Products = molecules resulting from reaction
6H2O + 6CO2 → C6H12O6 + 6O2
reactants products
26
• Extent of chemical reaction influenced by1. Temperature2. Concentration of reactants and products3. Catalysts
• Many reactions are reversible
27
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28
Water
• Life is inextricably tied to water
• Single most outstanding chemical property of water is its ability to form hydrogen bonds– Weak chemical associations that form
between the partially negative O atoms and the partially positive H atoms of two water molecules Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
a. Solid b. Liquid c. Gas
a: © Glen Allison/Getty Images RF; b: © PhotoLink/Getty Images RF; c: © Jeff Vanuga/Corbis
29
Polarity of water
• Within a water molecule, the bonds between oxygen and hydrogen are highly polar– Oxygen is much more electronegative than
Hydrogen
• Partial electrical charges develop– Oxygen is partially negative δ+
– Hydrogen is partially positive δ–
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+
–
+
+
a. b.
c.
Bohr Model
Space-Filling Model
8p
Ball-and-Stick Model
8n
–
+
+
–
–
+
+
–
104.5
+
31
Hydrogen bonds
• Cohesion – polarity of water allows water molecules to be attracted to one another
• Attraction produces hydrogen bonds• Each individual bond is weak and
transitory• Cumulative effects are enormous• Responsible for many of water’s important
physical properties
32
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+
a.
b.
Water moleculeHydrogen atom
–Hydrogen bond
Oxygen atom
Hydrogen atom
Hydrogen bond+
–
An organic molecule
Oxygen atom
• Cohesion – water molecules stick to other water molecules by hydrogen bonding
• Adhesion – water molecules stick to other polar molecules by hydrogen bonding
33
34
Properties of water
1. Water has a high specific heat– A large amount of energy is required to
change the temperature of water
2. Water has a high heat of vaporization– The evaporation of water from a surface
causes cooling of that surface
3. Solid water is less dense than liquid water
– Bodies of water freeze from the top down
35
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Hydration shells
Water molecules
Salt crystal
Cl–
Na+
Cl–
+
+
+
+
+
–
–
–
–
–
Na+
36
4. Water is a good solvent– Water dissolves polar molecules and ions
5. Water organizes nonpolar molecules– Hydrophilic “water-loving”– Hydrophobic “water-fearing”– Water causes hydrophobic molecules to
aggregate or assume specific shapes
6. Water can form ions
H2O OH– + H+
hydroxide ionhydrogen ion
37
Acids and bases
• Pure water– [H+] of 10–7 mol/L– Considered to be neutral– Neither acidic nor basic
• pH is the negative logarithm of hydrogen ion concentration of solution
Acids and Bases - pH
• Acid– Any substance that dissociates in water to
increase the [H+] (and lower the pH)– The stronger an acid is, the more hydrogen
ions it produces and the lower its pH
• Base– Substance that combines with H+ dissolved in
water, and thus lowers the [H+]
38
39
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Hydrogen IonConcentration [H+] Examples of Solutions
Stomach acid, lemon juice
pH Value
Hydrochloric acid
Acidic
Basic
10–2
10–3 Vinegar, cola, beer
10–4 Tomatoes
10–5 Black coffee
10–6 Urine
10–7 Pure water
10–8 Seawater
10–9 Baking soda
10–10 Great Salt Lake
10–11 Household ammonia
10–12
Household bleach10–13
10–14 Sodium hydroxide
0
1
2
3
4
5
6
8
9
10
11
12
13
14
100
10–1
Buffers
• Substance that resists changes in pH
• Act by – Releasing hydrogen ions when a base is
added– Absorbing hydrogen ions when acid is added
• Overall effect of keeping [H+] relatively constant
40
41
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1X0 3X 4X 5X2X
pH
9
876543210
Amount of base added
Buffering range
42
• Most biological buffers consist of a pair of molecules, one an acid and one a base
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+Bicarbonate
ion(HCO3
–)+
++
–
+
Water(H2O)
Carbondioxide(CO2)
Carbonicacid
(H2CO3)
Hydrogenion(H+)