1. Chapter 2 Life, Chemistry, and Water

2. Why It Matters. . .
Living organisms are collections of atoms and molecules linked together by chemical bonds
The laws of chemistry govern both living and nonliving things
Understanding the relationship between the structure of chemical substances and their behavior is the first step in biology

3. 2.1 The Organization of Matter
An element is a pure substance that cannot be broken down into simpler substances by ordinary chemical or physical techniques
All matter (anything that occupies space and has mass) is composed of elements and combinations of elements

4. 11 Key Elements Four elements: carbon, hydrogen, oxygen, and nitrogen
make up more than 96% of the weight of living organisms,
Chnops the cyclops
Seven elements (calcium, phosphorus, potassium, sulfur, sodium, chlorine, and magnesium) make up nearly 4%

5. Seawater Human Pumpkin Earth’s crust Oxygen 88.3 Hydrogen 11.0
Chlorine
Sodium
Magnesium
Sulfur
Potassium
Calcium
Carbon
Silicon
Nitrogen
Strontium
Oxygen
Carbon
Hydrogen
Nitrogen
Calcium
Phosphorus
Potassium
Sulfur
Sodium
Chlorine Magnesium
Iron
Iodine
Oxygen
Hydrogen
Carbon
Potassium
Nitrogen
Phosphorus
Calcium
Magnesium
Iron
Sodium
Zinc
Copper
Oxygen
Silicon Aluminum
Iron
Calcium
Sodium
Potassium
Magnesium
Other
elements
FIGURE 2.1 The proportions by mass of different elements in seawater, the human body, a fruit, and Earth’s crust.
Trace elements in humans include boron, chromium, cobalt, copper, fluorine, iodine, iron, manganese, molybdenum, selenium, tin, vanadium, and zinc
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6. Atoms and Molecules
Elements are composed of atoms – the smallest units that retain the chemical and physical properties of an element
Atoms combine chemically in fixed numbers and ratios to form the molecules of living and nonliving matter
Molecular names are written as a chemical formula, using standard symbols for elements and subscripts for the number of atoms of each element (such as CO2)

7. Compounds
Molecules whose component atoms are different (such as carbon dioxide) are called compounds
The chemical and physical properties of compounds are typically distinct from those of their atoms or elements
Example: Hydrogen (H) and oxygen (O) – the elements that form liquid water (H2O) – are both highly flammable gases

8. 2.2 Atomic Structure Each element consists of one type of atom
Each atom consists of an atomic nucleus surrounded by fast-moving, negatively-charged electrons
Atomic nuclei contain positively charged protons and uncharged neutrons

9. Atomic Structure
FIGURE 2.2 Atomic structure. The nucleus of an atom contains one or
more positively charged protons and, except for the most common form of
hydrogen, a similar number of uncharged neutrons. Fast-moving negatively
charged electrons, in numbers equal to the protons, surround the nucleus.
The most common form of hydrogen, the simplest atom, has a single proton
in its nucleus and a single electron (A). Carbon, a more complex atom, has a
nucleus surrounded by electrons at two levels (B). The electrons in the outer
level follow more complex pathways than shown here.

10. Protons and Neutrons Atoms are identified by their atomic number
Number of protons in the nucleus that does not vary
All atoms except hydrogen also contain at least one neutron
Atoms are assigned a mass number based on the total number of protons and neutrons in the atomic nucleus
Electrons contribute no significant mass

11. Isotopes
Isotopes are distinct forms of atoms of an element with the same number of protons but different numbers of neutrons
Organisms can use any hydrogen or carbon isotope, for example, without a change in their chemical reactions
Carbon 12 and life and carbon dating done in semester 2. Mention this

12. Atomic Nuclei of Isotopes
FIGURE 2.3 The atomic nuclei of hydrogen and carbon isotopes.
Note that isotopes of an atom have the same atomic number but different
mass numbers. Students do not need to know these specific examples, but it is nice to give an example to show what is an isotope. It comes up in GBIO 153 and carbon dating.

13. Unstable Nuclei
Nuclei of some isotopes are unstable and break down (decay) giving off particles of matter and energy (radioactivity)
Decay transforms the unstable, radioactive isotope (radioisotope) into an atom of another element (e.g., 14C to nitrogen)
Radioactive decay continues at a steady, clocklike rate

14. Isotopes in Research
Radiometric dating uses radioisotope decay to estimate the age of organic material, rocks, or fossils that contain them

15. Electron Orbitals
In an atom, the number of electrons equals the number of protons in the nucleus – making the atom electrically neutral
The region of space where an electron is found most frequently is its orbital
electrons may pass from one orbital to another within an atom, or pass completely to another atom forming a bond between the two atoms

16. Electron Shells
Within an atom, electrons are found in regions of space called energy levels or shells
The shell nearest to the nucleus may be occupied by a maximum of two electrons in a single, spherical orbital
The 2nd and 3rd energy levels can hold a maximum of eight electrons.

17. Electron Shells (cont’d.)
Energy levels are filled in order from inner to outer shells
The total number of electrons in the orbitals matches the number of protons in the nucleus
The electrons in the outer energy levels have more energy than the inner energy levels

18. Elements and Electron Shells
FIGURE 2.5 The atoms with electrons distributed in one, two, or three energy levels.
The atomic number of each element (shown in boldface in each panel) is equivalent to the number of protons in its nucleus.
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19. Electrons Determine Chemical Activity
The electrons in an atom’s outermost energy level are its valence electrons
Atoms with a completely filled outermost energy level are nonreactive, or inert (helium, neon, argon)
Atoms in which the outermost energy level is not completely filled with electrons are chemically reactive (hydrogen)
Chemical bonds/ chemical activity fill the valence energy level.

20. Electrons Determine Chemical Activity (cont'd.)
The relative tendency of an atom to gain, lose, or share valence electrons underlies the chemical bonds and forces that hold molecules together
Octet rule- refers to 8 valance electrons required to fill the 2nd or 3rd energy level to make an atom stable. An atom will gain, lose, or share valence electrons to become stable.

21. Electrons Determine Chemical Activity (cont'd.)
Atoms with outer shells that contain electrons near the stable numbers tend to gain or lose electrons (e.g., sodium, chlorine) (Ionic bonds)

22. Electrons Determine Chemical Activity (cont'd.)
Atoms that differ from the stable configuration by more than one or two electrons tend to share electrons (covalent bonds) in joint orbitals with other atoms (e.g., oxygen, carbon)

23. 2.3 Chemical Bonds and Chemical Reactions
Atoms of reactive elements tend to combine into molecules by forming chemical bonds
The three most important chemical bonds in biological molecules:
ionic bonds
covalent bonds
hydrogen bonds
Chemical reactions either form new chemical bonds or break old ones
Van der Waals forces come up in the MindTap Assignment, so you may want to mention them, but they will not be tested on them.

24. Ionic Bonds
Ionic bonds result from electrical (+ -) attractions between atoms that gain or lose valence electrons completely (ions)
Ionic bonds are strong until dissolved in water where they break.
A positively charged ion (one that has lost an electron), such as Na+, is called a cation
A negatively charged ion (one that has gained an electron), such as Cl-, is called an anion

25. Electron loss Electron gain A. Sodium atom 11 e– 11 p+ Chlorine atomCl
FIGURE 2.6 Formation of an ionic bond.
Sodium, with one electron in its outermost energy level, readily loses that electron to attain a stable state in which its second energy level, with eight electrons, becomes the outer level. Chlorine, with seven electrons in its outer energy level, readily gains an electron to attain the stable number of eight. The transfer creates the ions Na+ and Cl-. The combination forms sodium chloride (NaCl), common table salt.
Sodium
ion
10 e–
11 p+
Chlorine
ion
18 e–
17 p+
Na+
Cl–

26. Formation of an Ionic Bond
B. Combination of sodium and chlorine in sodium chloride (table salt)
Cl–
Na+
FIGURE 2.6 Formation of an ionic bond.
Sodium, with one electron in its outermost energy level, readily loses that electron to attain a stable state in which its second energy level, with eight electrons, becomes the outer level. Chlorine, with seven electrons in its outer energy level, readily gains an electron to attain the stable number of eight. The transfer creates the ions Na+ and Cl-. The combination forms sodium chloride (NaCl), common table salt.
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27. Covalent Bonds
Covalent bonds form when atoms share a pair of valence electrons
In molecular diagrams a covalent bond is designated by a pair of dots or a single line that represents a pair of shared electrons
Example: H2 is represented as H:H or H—H

28. Polarity
The higher the electronegativity of an atom, the more strongly it attracts shared electrons
Covalent bonds differ in the sharing of valence electrons, depending on the difference in electronegativity between the bonded atoms

29. Polarity (cont'd.)
In a nonpolar covalent bond, electrons are shared equally, no charges
H2 and O2
CH4, C-H bonds form nonpolar molecules

30. Polarity (cont'd.)
In a polar covalent bond, electrons are shared unequally
The atom that attracts the electrons more strongly carries a partial negative charge, δ-
The atom deprived of electrons carries a partial positive charge, δ+
The whole molecule is polar, because one end is partially positive and the other end is partially negative

31. Water, a Polar Molecule
In water, an oxygen atom forms polar covalent bonds with two hydrogen atoms
The electrons are attracted much more strongly to the oxygen nucleus than to the hydrogen nuclei
The oxygen atom is located on one side (δ-) and hydrogen atoms on the other (δ+)

32. Polarity in Water
FIGURE 2.8 Polarity in the water molecule, created by unequal
electron sharing between the two hydrogen atoms and the oxygen
atom and the asymmetric shape of the molecule. The unequal
electron sharing gives the hydrogen end of the molecule a partial positive
charge, d1 (“delta plus”), and the oxygen end of the molecule a partial
negative charge, d2 (“delta minus”). Regions of deepest color indicate the
most frequent locations of the shared electrons. The orbitals occupied by
the electrons are more complex than the spherical forms shown here.

33. Polar Molecules Associate and Exclude Nonpolar
Polar molecules attract other polar molecules and charged ions and molecules, forming polar associations that tend to exclude nonpolar molecules
Polar molecules that associate readily with water are hydrophilic (“water loving”)

34. Polar Molecules Associate and Exclude Nonpolar (cont’d.)
Excluded nonpolar molecules tend to clump together in nonpolar associations which reduce the surface area exposed to the surrounding polar environment
Nonpolar substances that are excluded by water and other polar molecules are hydrophobic (“water fearing”)

35. Hydrogen Bonds
Hydrogen bonds are attractions between partially positive hydrogen atoms (sharing electrons unequally with oxygen, nitrogen, or sulfur) and partially negative atoms sharing in a different covalent bond

36. Hydrogen Bonds (cont'd.)
Individual hydrogen bonds are weak, but when numerous, hydrogen bonds are collectively strong
Hydrogen bonds are found in large biological molecules such as proteins and DNA
Hydrogen bonds between water molecules are responsible for many of the properties that make water important to life

37. Hydrogen Bonds Hydrogen bond FIGURE 2.9 Hydrogen bonds.
(A) A hydrogen bond (dotted line) between the hydrogen of an —OH group and a nearby nitrogen atom, which also shares electrons unequally with another hydrogen. Regions of deepest blue indicate the most likely locations of electrons. (B) Multiple hydrogen bonds stabilize the backbone chain of a protein molecule into a spiral called the alpha helix. The spheres labeled R represent chemical groups of different kinds.

38. Stabilizing Effect of Hydrogen Bonds
FIGURE 2.9 Hydrogen bonds.
(A) A hydrogen bond (dotted line) between the hydrogen of an —OH group and a nearby nitrogen atom, which also shares electrons unequally with another hydrogen. Regions of deepest blue indicate the most likely locations of electrons. (B) Multiple hydrogen bonds stabilize the backbone chain of a protein molecule into a spiral called the alpha helix. The spheres labeled R represent chemical groups of different kinds.

39. 2.4 Hydrogen Bonds and the Properties of Water
Hydrogen bonds between water molecules produce a water lattice that affects properties of density, heat absorption, cohesion, and surface tension
Polarity of water molecules contributes to formation of distinct polar and nonpolar environments critical to cell organization
Water is a solvent for charged or polar molecules
Water molecules separate into hydrogen and hydroxyl ions.

40. A Lattice of Hydrogen Bonds
Liquid water forms a water lattice
Each water molecule makes multiple H bonds, which constantly break and reform
An ice lattice is a rigid, crystalline structure
the water molecules are farther apart than the water lattice
Ice is less dense than liquid water, an unusual property that makes ice float

41. Hydrogen Bonds and Water
A. Hydrogen-bond lattice of liquid water
B. Hydrogen-bond lattice of ice
FIGURE 2.11 Hydrogen bonds and water.
In liquid water, hydrogen bonds (dotted lines) between water molecules produce a water lattice. The hydrogen bonds form and break rapidly, allowing the molecules to slip past each other easily.
(B) In ice, water molecules are fixed into a rigid lattice.
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42. Water and Temperature
The hydrogen-bond lattice of liquid water stabilizes water as it is heated
Water remains liquid in a wide temperature range (0oC to 100oC)
Water has a relatively high specific heat
It can absorb or release relatively large quantities of heat energy without undergoing extreme changes in temperature

43. Specific Heat and Calories
Amount of heat energy required to increase the temperature of a given quantity of water
Measured in calories
Calorie (or small calorie)
Heat energy required to raise 1 g of water by 1°C
Calorie (with a capital C)
A kilocalorie (kcal) or 1,000 calories

44. Heat of Vaporization
A large amount of heat (586 calories per gram) must be added to give water molecules enough energy of motion to break loose from liquid water and form a gas
This required heat, known as heat of vaporization, allows humans and many other organisms to cool off when hot

45. Cohesion and Adhesion Cohesion Adhesion
Tendency of water molecules to “stick” to each other, due to the hydrogen-bond lattice
Adhesion
Tendency of water molecules to “stick” to the walls of tubes by forming hydrogen bonds with charged and polar groups in molecules that form the walls

46. Surface Tension
Water molecules at surfaces facing air form additional hydrogen bonds with water molecules beside and below them, but not on the sides that face the air
This bonding places the surface water molecules under tension (surface tension) making them more resistant to separation than the underlying water molecules
Surface tension causes water to form water droplets, and can support small insects and other objects

47. Surface Tension in Water
FIGURE Surface tension in water.
(A) Unbalanced hydrogen bonding places water molecules under lateral tension where a water surface faces the air.
(B)A raft spider (Dolomedes fimbriatus) is supported by the surface tension of water.

48. Polar and Nonpolar Environments
The water lattice resists invasion by other molecules unless the molecule contains polar or charged regions that can form competing attractions with water molecules
Nonpolar molecules are excluded from the water lattice, forcing them to form nonpolar associations that expose the least surface area to the surrounding water
Distinct polar and nonpolar environments created by water are critical to the organization of cells

49. Water as a Solvent
Water (solvent) surrounds the dissolved substance (solute) and prevents the polar molecules or ions from reassociating (e.g., sodium and chloride)

50. Water Forming a Hydration Layer
FIGURE Water molecules forming a hydration layer around Na1
and Cl2 ions, which promotes their separation and entry into solution.