BIOL 1020 - CHAPTER 2 LECTURE NOTES

Chapter 2: You must understand chemistry to understand life

1. Describe the difference between the terms “element” and “atom”. What are chemical symbols, and what is the periodic table?

2. Draw a model of a neutral atom with atomic number 6 and atomic mass 12, and compare with your groupmates.

3. Draw a model of a neutral atom with atomic number 6 and atomic mass 14, and compare with your groupmates. Then compare with the previous atom that you drew and discuss isotopes.

4. What are electron orbitals? What is the valence shell? How does the valence shell relate to chemical reactivity of an atom?

5. As a group, draw a Venn diagram for the following terms: molecule, compound. Then place the following in the diagram: O2, NaCl, H2O

6. Discuss moles, atomic mass units, and Avogadro’s number. Why use moles instead of just mass? Discuss the water and glucose problem.

7. What is a covalent bond? What do polar and nonpolar mean for covalent bonds? Give an example of each.

8. What are ions? What are cations and anions? What is an ionic bond? Give an example.

9. What are hydrogen bonds? Draw an example.

10. Discuss the chemical equations in the notes and the terms there (reactant, product, equilibrium).

11. What is a redox reaction, and how does it relate to movement of electrons and movement of energy? What gets oxidized/reduced in the following: making NaCl; rusting iron What gains/loses energy in each case?

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BIOL 1020 - CHAPTER 2 LECTURE NOTESOverview: In many ways, life can be viewed as a complicated chemical reaction. Modern models of how life works at all levels typically have at least some aspect of chemistry as a major component or underpinning.

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BIOL 1020 - CHAPTER 2 LECTURE NOTESI. Elements and Atoms

A. elements – substances that cannot be further broken down into other substances (at least by ordinary chemical reactions)

1. every element has a chemical symbol (H for hydrogen, O for oxygen, etc.); this is most familiar from the periodic table

2. there are 92 naturally occurring elements, from hydrogen up to uranium

4 elements (O, C, H, N) make about 96% of the mass of most living things

8 others are consistently present in small amounts in living things (Ca, P, K, S, Na, Mg, Cl, Fe)

several others are typically found only in trace amounts (trace elements); these tend to vary considerably in amount

and even presence depending on the type of organism

B. an atom is the smallest unit of an element that still retains the properties of that element

C. atoms consist of subatomic particles

1. electron - contributes no significant mass to the atom, but carries a (-1) electrical charge

2. proton - contributes a mass of approximately 1 mass unit, and carries a (+1) electrical charge

3. neutron - contributes a mass of approximately 1 mass unit, and carries no net electrical charge

4. protons and neutrons are found in the nucleus (center) of an atom

5. elements differ from each other because they contain different numbers of protons (all hydrogen atoms contain 1 proton,

all carbon atoms contain 6 protons, all oxygen atoms contain 8 protons, etc.)

atomic number = number of protons in the nucleus

the periodic table has elements arranged largely according to atomic number

6. protons + neutrons determine atomic mass

each contribute ~1 atomic mass unit (amu, or Dalton)

atoms that have the same number of protons but have different numbers of neutrons (therefore different masses) are

referred to as isotopes

D. atomic nuclei can undergo changes (decay)

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BIOL 1020 - CHAPTER 2 LECTURE NOTES1. some elements are more stable than others

2. some isotopes are more stable than others (most unstable = radioisotopes)

3. decay rates are statistical averages; used for measuring time passage in many areas of science (carbon dating, etc.)

4. the radiation emitted upon decay (alpha, beta, and/or gamma) can be used as a tool for experiments; can also be used

medically; has other uses and dangers (nuclear power, nuclear bombs, radiation poisoning, etc.)

5. radiation can cause mutations in DNA, can interfere with cell division

E. electrons occupy orbitals surrounding the nucleus and move at the speed of light

1. because ATOMS are electrically neutral the number of electrons an atom has always equals the number of protons

2. electrons can exist at different energy levels, which correspond to orbitals

the further away an orbital carries an electron from the nucleus, the higher the energy level of the electron

electrons with similar energies make up an electron shell

3. the outer electron(s) are known as the valence electron(s); collectively, they occupy the valence shell

4. the chemical properties of an atom are largely determined by the valence electrons

F. the science of chemistry mostly involves study of how electrons move about the nucleus, store energy, and determine

chemical properties of substances as a result

II. Describing Atomic Combinations

A. atoms combine to form molecules and compounds

1. molecule – two or more atoms held together by covalent bonds (defined later)

may be composed of one or more elements (examples: O2, H2O)

not all substances are molecular (NaCl, table salt, isn’t)

if a substance is molecular, then an individual molecule is the smallest unit of the substance that exhibits the

properties of the substance

thus, a molecule differs in its physical and chemical properties from the elements that make it up

2. compound - a specific combination of two or more different elements chemically combined in a fixed ratio

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BIOL 1020 - CHAPTER 2 LECTURE NOTES compounds have unique physical and chemical properties that differ from those of the elements used to make it

some compounds are held together by covalent bonds and are therefore molecular; some are held together by ionic

bonds (defined later)

B. chemists use two types of formulas to describe substances

1. chemical formula - a shorthand formula showing the number of atoms of each element present in a molecule

often called molecular formula if a molecule is involved; examples: H2O, CO2, O2, C6H12O6

follows simplest ratio for ionic substances (NaCl, etc.)

2. structural formula - shows the arrangement of atoms in a molecule

examples:

water H─O─H

carbon dioxide O═C═O

molecular oxygen O═O

C. the number of units of a substance are described using the mole

1. molecular mass is the sum of the atomic masses of the atoms in the molecule

2. since the actual mass of an atom is extremely small, it is convenient in the real world to work with a large number of

atoms at the same time

3. The amount of a substance that in grams has the same number as the atomic mass is a mole

4. Thus, water has molecular mass 1+1+16 = 18; a mole of water has a mass of 18 g

5. The mole is simply a conversion factor from the small scale of atomic mass units to the more familiar gram scale

the factor represents the number of units (molecules or atoms) in a mole

this factor, called Avogadro’s number, is 6.02 x 1023 atoms or molecules

III. Chemical Bonds Hold Molecules Together and Store Energy

A. recall that electrons in the outermost shell of an atom (valence electrons) determine the chemical behavior of the atom, i.e.

what type and how many chemical bonds it can readily form

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BIOL 1020 - CHAPTER 2 LECTURE NOTESB. most atoms in biological systems seek to have 8 electrons in their outermost shell (hydrogen seeks to have 0 or 2 electrons in

its outermost shell)

C. since atoms have the same number of electrons as protons, they meet this need to have a full valence shell by sharing, giving

up, or acquiring electrons from other atoms; this forms chemical bonds

1. a chemical bond is a reduced energy state

2. bond energy is the amount of energy required to break a particular chemical bond

D. there are two principle types of strong chemical bonds

1. covalent bonds - electrons are shared between two atoms

2. ionic bonds - one atom completely gives up an electron to another atom

E. covalent bonds

1. result in filled valence shells

2. electrons are shared in pairs

3. a single electron pair shared = a single covalent bond

4. double and triple covalent bonds are also possible

5. carbon forms 4 covalent bonds

6. covalent bonds give molecules definite shapes

the shared atomic orbitals require definite spatial arrangements that depend on the atoms involved in the bond

7. covalent bonds can be nonpolar (equal sharing of electrons) or polar (unequal sharing of electrons)

polar bonds result if one nucleus holds a stronger attraction on the electron pair

molecules with polar bonds (polar molecules) have regions with partial charges

F. ionic bonds

1. when an atom gains or gives up one or more electrons, it is called an ion

cations - ions that have lost one or more electrons; have a positive charge

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BIOL 1020 - CHAPTER 2 LECTURE NOTES anions - ions that have gained one or more electrons; have a negative charge

the suffix –ide indicates an anion

2. polyatomic ions can also form

covalently bound atoms that lose or gain electrons or protons

only polyatomic ions can lose or gain protons

polyatomic cations = positive charge; polyatomic anions = negative charge

3. an ionic bond is formed by the attraction between a cation and an anion

4. an ionic compound is a substance held together by ionic bonds

ionic compounds dissociate into individual ions when dissolved in a polar substance, such as water

hydration – surrounding the ions with the ends water molecules with the opposite (partial) charge

G. hydrogen bonds

1. weak interactions involving partially (+) charged hydrogen atoms

2. the interaction is with another atom with a partial (-) charge

3. can be within the same (large) molecule, or between molecules

4. hydrogen bonds are common and important in living things

water forms hydrogen bonds

because they are weak, hydrogen bonds are relatively easy to manipulate

collectively, hydrogen bonds can be very strong – they hold together the two strands of DNA, for example

H. In aqueous systems (such as living organisms), the typical relative bond strengths are:

covalent bond > ionic bond > hydrogen bond

IV. Chemical Equations Describe Chemical Reactions

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BIOL 1020 - CHAPTER 2 LECTURE NOTESA. Reactants are written on the left

B. Products are written on the right

C. an arrow ( ) is used to show the direction the reaction proceeds

C6H12O6 + 6 O2 6 CO2 + 6 H2O + Energy

D. double arrows of equal lengths ( ) indicate equilibrium reactions (reactions proceeding simultaneously at equal

rates in both directions)

N2 + 3 H2 2 NH3

E. Sometimes, different lengths of double arrows are used to indicate which direction is favored

CO2 + H2O H2CO3

V. Oxidation-Reduction Reactions (redox reactions) Are Common in Biological Systems

A. oxidation is a chemical process in which an atom, molecule, or ion loses an electron(s)

B. reduction is the opposite – an electron is gained (charge is reduced)

C. oxidation and reduction are always paired (hence redox reactions)

D. example: rusting

1. when iron rusts, iron oxide is formed by the oxidation of iron; this can be described by a chemical reaction as shown

below:

4 Fe + 3 O2 2 Fe2O3

2. during the process iron atoms (Fe) become iron ions (Fe3+):

4 Fe 4 Fe3+ + 12 e-

3. therefore, we can say that iron atoms were oxidized to produce iron ions above

4. on the flip side, the oxygen atoms gain electrons; we can say that the oxygen is reduced in the reaction:

3 O2 + 12 e- 6 O2-

E. oxygen is the most common oxidizing agent (hence the general term oxidation)

F. in biological systems, typically molecules are oxidized and reduced

1. very important in many processes such as photosynthesis, respiration

2. electrons are less easily lost from molecules than from atoms

molecules typically will lose the equivalent of a complete hydrogen atom when oxidized

this means that both a proton and an electron are removed from the oxidized molecule and may be added to the

reduced molecule

thus, counting charge changes is not sufficient for analyzing redox reactions – look for movement of electrons in

redox reactions involving biological molecules

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