Lab Activity 2
mattymatt121
Mason_UB3e_ch02_PPT_final_Accessible.pptxChapter 2 Lecture Outline
Understanding Biology
THIRD EDITION
Kenneth A. Mason
Tod Duncan
Jonathan B. Losos
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The Nature of Molecules and the Properties of Water
Chapter 2
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All Matter is Composed of Atoms
Matter has mass and occupies space
All matter is composed of atoms
Atoms are composed of subatomic particles
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3
Atomic Structure
Atoms are composed of three types of subatomic particles
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
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Figure 2.2
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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
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Atomic mass
Mass or weight?
Mass – refers to amount of substance
Weight – refers to the force gravity exerts on a substance
Sum of protons and neutrons is the atom’s atomic mass
Each proton and neutron has a mass of approximately 1 Dalton
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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 charge
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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
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Figure 2.3
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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
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Figure 2.4
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Atomic 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
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Figure 2.5
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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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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
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Periodic Table of the Elements
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Figure 2.6b
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,
Organic molecules contain primarily CHON
Some trace elements are very important
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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
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Ionic bonds
Formed by the attraction of oppositely charged ions by electrostatic force
Ions form when the atom has a gain or loss of electrons
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
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Figure 2.8
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Covalent bonds 1
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
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Covalent bonds 2
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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
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Hydrogen bonds
Electropositive hydrogen from one polar molecule is attracted to an electronegative atom that is often oxygen
Attraction produces hydrogen bonds
Each individual bond is weak and transitory
Cumulative effects are enormous
Responsible for many of water’s important physical properties
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Van de Waals Attraction
Weak bond
Non-directional attractive force called Van der Waals forces
Form when two atoms are very close to one another
Antibodies recognize the shape of an invading organism with this bond
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Chemical reactions 1
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
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Chemical reactions 2
Extent of chemical reaction influenced by
Temperature
Concentration of reactants and products
Catalysts
Many reactions are reversible
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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
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Figure 2.9
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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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Figure 2.10
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Figure 2.11
Surface tension of water
Cohesion – water molecules stick to other water molecules by hydrogen bonding
Surface tension due to hydrogen bonds
© Hermann Elsenbeiss/National Audubon Society Collection/Science Source.
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Figure 2.12
Adhesion – water molecules stick to other polar molecules by hydrogen bonding
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Properties of water 1
TABLE 2.3 The Properties of Water
| Property | Explanation | Example of Benefit to Life |
| Cohesion/Adhesion | Hydrogen bonds cause water molecules to be attracted to other polar or charged species. | Leaves pull water upward from the roots; seeds swell and germinate. |
| High specific heat | Hydrogen bonds absorb heat when they break and release heat when they form, minimizing temperature changes. | Water stabilizes the temperature of organisms and the environment. |
| High heat of vaporization | Many hydrogen bonds must be broken for water to evaporate. | Evaporation of water cools body surfaces. |
| Lower density of ice | Water molecules in an ice crystal are spaced relatively far apart because of hydrogen bonding. | Because ice is less dense than water, lakes do not freeze solid, allowing fish and other life in lakes to survive the winter. |
| Solubility | Polar water molecules are attracted to ions and polar compounds, making these compounds soluble. | Many kinds of molecules can move freely in cells, permitting a diverse array of chemical reactions. |
| Hydrophobic exclusion | Water repels hydrophobic compounds, forcing them to associate together. | Biological membranes have bilayer structure with hydrophobic interior. |
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Properties of water 2
Water has a high specific heat
A large amount of energy is required to change the temperature of water
Water has a high heat of vaporization
The evaporation of water from a surface causes cooling of that surface
Solid water is less dense than liquid water
Bodies of water freeze from the top down
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Figure 2.13
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Properties of water 3
Water is a good solvent
Water dissolves polar molecules and ions
Water organizes nonpolar molecules
Hydrophilic “water-loving”
Hydrophobic “water-fearing”
Water causes hydrophobic molecules to aggregate or assume specific shapes
Water can form ions
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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
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Acids and Bases - pH
Acid
Any substance that dissociates in water to increase the [H+] (and lowers 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+] (and raises the pH)
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Figure 2.14
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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
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Figure 2.15
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Biological buffers
Most biological buffers consist of a pair of molecules, one an acid and one a base
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Accessibility Content: Text Alternatives for Images
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Figure 2.2 - Text Alternative
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Hydrogen has one positively charged proton in the nucleus and one negatively charged electron in an orbital. Oxygen has 8 protons and 8 neutrons (with no charge) in the nucleus and 8 electrons, two in an inner orbital and 6 in the outer orbital.
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Figure 2.3 - Text Alternative
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The three isotopes of carbon differ in their number of neutrons, Carbon-12 has 6 neutrons, Carbon-13 has 7 neutrons, Carbon-14 has 8 neutrons.
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Figure 2.4 - Text Alternative
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Neon has 10 electrons, 2 in the inner shell and 8 in the outer shell. The 2 electrons in the inner shell are in a spherical 1s orbital. Two of the electrons in the outer shell are in a spherical 2s orbital. The remaining 6 electrons in the outer shell are in pairs in three dumbbell shaped p orbitals.
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Figure 2.6b - Text Alternative
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Carbon, oxygen, hydrogen, nitrogen, sodium, chlorine, calcium, phosphorous, potassium, sulfur, iron, magnesium
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Figure 2.8 - Text Alternative
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When sodium gives an electron to chlorine, both have full outer shells. The positive charge on the sodium ion is attracted to the negative charge on the chloride ion, and this attraction forms a sodium chloride salt crystal.
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Covalent bonds 2 - Text Alternative
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Hydrogen gas has two hydrogen atoms sharing a pair of electrons to form one covalent bond. Oxygen gas has two oxygen atoms sharing two pairs of electrons to form two covalent bonds. Nitrogen gas has two nitrogen atoms sharing three pairs of electrons to form three covalent bonds.
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Figure 2.9 - Text Alternative
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The oxygen has more pull on the electrons, giving it a partial negative charge and the hydrogen atoms a partial positive charge.
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Figure 2.10 - Text Alternative
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Two water molecules have a hydrogen bond between the positive charge on the hydrogen atom of one water molecule and the oxygen atom on the second water molecule. An organic molecule with an OH group can also form a hydrogen bond with water through its hydrogen atom and the oxygen atom of water.
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Figure 2.13 - Text Alternative
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A salt crystal dissolves in water when the Na+ ions interact with partial negative charges on water and Cl- ions interact with partial positive charges on water.
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Figure 2.14 - Text Alternative
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A hydrogen ion concentration of 10 -1 M has a pH of 1 (acidic) down to a concentration of 10 -14 M has a pH of 14 (basic). A list of the pH value of different solutions: hydrochloric acid pH 1, stomach acid pH 2, vinegar pH 3, tomatoes pH 4, coffee pH 5, urine pH 6, water pH 7, sea water pH 8, baking soda pH 9, great salt lake pH 10, ammonia pH 11, bleach pH 13, sodium hydroxide pH 14.
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Figure 2.15 - Text Alternative
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Increasing base is on the x-axis and pH on the y-axis. As base is added the pH increases rapidly at first, then the curve flattens out. This is the buffering range as adding base does not affect pH much. At the end of the graph pH again increases rapidly as base is added.
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Biological buffers - Text Alternative
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Carbonic acid breaks down to bicarbonate ion (HCO3-) and a hydrogen ion (H+).
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