chapter 2 - the chemical basis of life

matter

matter

anything that occupies space and has mass

• found on earth in three physical states: solid, liquid, and gas

  *living organisms are made of matter

element

a substance that cannot be broken down into other substances by ordinary chemical means

• each element has a symbol made up of the first letter or two of its english, latin, german, name (O - oxygen)

  *types of matter are all composed of elements

compound

a substance consisting of two or more different elements combined in a fixed ratio

• compounds have different characteristics than the elements that make up the compound

  • ex. H and O (gases) -→ H20 (water)

  *key idea of emergent properties

major elements essential for life

*first six elements make up 99% of the human body

1. oxygen (O): 65%

2. carbon (C): 18.5%

3. hydrogen (H): 9.6%

4. nitrogen (N): 3.3%

5. calcium (Ca): 1.5%

6. phosphorus (P): 1.0%

7. potassium (K): 0.4%

8. sulfur (S): 0.3%

9. sodium (Na): 0.2%

10. chlorine (Cl): 0.2%

11. magnesium (Mg): 0.1%

the four major elements of biological molecules

1. oxygen (O)

2. carbon (C)

3. hydrogen (H)

4. nitrogen (N)

oxygen (O)

• atomic number: 8

  • when neutrally charged, it has 8 electrons (6 in its valence shell)

    *can make TWO bonds

• electronegativity: 3.44

  • highest of the four main elements

carbon (C)

• atomic number: 6

  • when neutrally charged, it has 6 electrons (4 in its valence shell)

    *can make FOUR bonds - vital in making organic compounds

• electronegativity: 2.55

  • similar to that of hydrogen (forms nonpolar covalent bonds with H)

hydrogen (H)

• atomic number: 1

  • when neutrally charged, it has 1 electron (1 in its valence shell)

    *can make ONE bond

• electronegativity: 2.2

  • similar to that of carbon (forms nonpolar covalent bonds with C)

nitrogen (N)

• atomic number: 7

  • when neutrally charged, it has 7 electrons (5 in its valence shell)

    *can make THREE bonds

• electronegativity: 3.04

elements in which are components for bone and teeth

• calcium (Ca)

• phosphorus (P)

most of elements in the remaining 1% of the body

*these elements are involved in functions such as nerve signalling and chemical reactions (redox)

• potassium (K)

• sulfur (S)

• sodium (Na)

• chlorine (Cl)

• magnesium (Mg)

trace elements

present elements in minute quantities, making up less than 0.01% of the human body

• baron, chromium, cobalt, copper, fluorine, iodine, iron, manganese, molybdenum, selenium, silicon, tin, vanadium, and zinc

  • some elements are required by some organisms but not all

    • iodine (I) is for organisms only with a vertebrate (a backbone)

  • ex. iron (Fe) is a trace element needed by all forms of life (vital for energy processing and transporting oxygen into the blood) — 0.004%

iodine as a trace element

iodine is an essential component of a hormone produced by the thyroid gland

• a deficiency in the diet causes the thyroid gland to grow to an abnormal size (condition is called a goiter)

• serious effects take place during fetal development and childhood

  • miscarriages, poor growth, mental impairment

• seafood, kelp, dairy, and leafy greens are great sources

  • inland regions are common for iodine deficiencies

    *universal consumption of iodization of all salt (iodized salt)

iron as a trace element

iron deficiency is the most common nutritional disorder with two billion people affected (mainly children and women)

• to avoid, food fortification, iron supplements, and diet diversification/improvement

  • however, too much iron could be fatal and could damage organs and increase risk of disease

fluoride as a trace element

fluoride, a form of fluorine, helps with maintaining strong bones and public health

• is added to our water (government controls this —> debates whether or not they should have the right too or not)

• fluoride in toothpaste helps maintain integrity of the tooth and prevent cavities

  • too much makes the teeth look brown, but significantly reduces the amount of cavities

atom

the smallest unit of matter that still retains the properties of an element

• each element has its own type of atom, which is different from the atoms of other elements

subatomic particles

parts of an atom

• proton

• electron

• neutron

proton

a subatomic particle with a single positive electrical charge (+)

• the amount of proton determines the element of the atom

electron

a subatomic particle with a single negative electrical charge (-)

• the distribution if electrons determines an atom’s chemical properties

  • only electrons are directly involved in the chemical activity of an atom

neutron

a subatomic particle that is electrically neutral (no charge)

• the amount of neutrons determines the stability of the atom and forms different isotopes.

nucleus

the atom’s central core

• made up of protons and neutrons

electron cloud

formed by negative charge by the two electrons surrounding the nucleus

• attraction between the electrons and protons holds the electrons near the nucleus

atomic number

how many protons an atom has

• the amount of protons an atom has determine the element

  *unless indicated, an atom has an equal number of protons and electrons (net charge = 0)

atomic mass

is approximately equal to its mass number - the sun and its protons and neutrons - in daltons (commonly an average of isotopes)

radioactive isotopes with basic research

• used carbon dioxide (CO2) containing the radio isotope, carbon-14, to study photosynthesis

  • allowed researchers to trace the sequence of molecules made by plants in the chemical route from CO2 to sugar

radioactive isotopes with medical diagnosis and treatment

• used to tag chemicals that accumulate in specific areas of the body (ex. phosphorus in bones)

  • scans and camericas can show where the radiation collects

• using radioactive iodine to kill the cancer in the thyroid as it usually accumulates there

  • PET (positron-emission technology) scan, can produce images of areas of the body with high metabolic activity (with using glucose/oxygen that is radioactive)

    *extremely helpful in diagnosing heart disorders, cancers, and aiding in brain research such as alzheimer’s

dangers of radioactive isotopes

• particles and energy thrown off my radioactive atoms can damage molecules and DNA

  • explosion of a nuclear reactor in chernobyl (ukraine - 1986)

  • radon (a radioactive gas) is the second leading cause of cancer

electron shells

locations in which an electron could be, each with a characteristic distance from the nucleus

• depending on an atom’s atomic number, an atom may have one, two, or more electron shells

  *the first shell holds TWO electrons, the second shell holds EIGHT electrons, the third shell holds EIGHTEEN electrons

orbitals

discrete volumes of space in which electrons are most likely to be found

• within each shell, electrons travel in different orbitals

• each orbital can hold a maximum of TWO electrons (ex. the 2nd shell has four orbitals and can hold up to eight electrons)

valence shell

the outermost electron shell of an atom

• the amount of electrons in that shell mostly determines the chemical property of an atom

  • atoms whose outer shells are not full tend to interact with other atoms in a way that enable them to complete or fill their valence shells (if the valence shell is full, they will rarely react with other atoms — usually classified as noble gases)

    *hydrogen, oxygen, nitrogen, and carbon react with other atoms to form bonds

valence/bonding capacity

the number of covalent bonds an atom can form depends on the number of electrons needed to fill its valence shell

• hydrogen needs 1

• oxygen needs 2

• nitrogen needs 3

• carbon needs 4

chemical bonds

attractions in which holds atoms together due to their interactions (giving up, accepting, or sharing electrons)

• ionic bonds: the transfer of electrons between atoms result in this bond

• covalent bonds: the sharing of electrons result in this bond

covalent bonds

joins atoms into molecules through electron sharing

molecule

consists of two or more atoms held together by covalent bonds

• the sharing of electrons is not always equal

electronegativity

a measure of attraction between atoms

• a low difference in electronegativity result in nonpolar covalent bonds

• a medium difference in electronegativity result in polar covalent bonds

• a large difference in electronegativity result in ionic bonds

polarity

refers to a separation of charges

• think of the north and south poles of earth, or a magnet

nonpolar covalent bonds

the electrons are shared EQUALLY between two atoms of the same of similar electronegativity

polar covalent bonds

the electrons are drawn more closely to the more electronegative element between the two atoms that differ in electronegativity

• as a result, the more electronegative atom carries a partially negative charge in the molecule and the other carries a partially positive charge in the molecule

molecular formula

shows the number of atoms of each element using symbols and subscripts

electron distribution diagram

shows how each atom completes its outer shell by sharing one or more pairs of electrons

structural formula

shows a molecule’s approximate shape and represents each covalent bond with a line

• one line represents a single bond

• two lines represents a double bond

space-filling model

uses a color-coded ball for each atom and comes closest to representing a molecule’s 3-D shape

ionic bonds

attractions between ions of opposite charge

• two atoms are so unequal in their attraction for electrons that the more electronegative atom strips an electron completely from another atom

  • ex. NaCl (table salt) — Na gives an electron to Cl to complete their valence shells in which their resulting charges form a bond

redox reaction

a chemical reaction in which involves a transfer of electrons between two atoms

• the transfer of an electron moves one unit of negative charge from one atom to another

  • reduction: when an atom loses an electron (becomes positively charged) — ex. Na

  • oxidization: when an atom gains an electron (becomes negatively charged) - ex. Cl

ions

an atom or molecule with an electrical charge resulting from a gain or a loss of one or more electrons

• when the attraction between two ions with opposite charges holds them together is an ionic bond

  • cation: positively charged ion

  • anion: negatively charged ion (names of anions often end in “-ide”)

salts

a synonym for an ionic compound

• affects of the environment on ionic bonds

  • when DRY, bonds are so strong you need a hammer and a chisel to break the crystal

  • when IN WATER, ions interact with polar water molecules and the salt dissolves

strong and weak bonds

• covalent bonds are strong, linking atoms to a cell’s molecules

• ionic bonds are weaker, but crucial to the functioning of a cell concerning bonds within and between molecules

• hydrogen bonds are an important week bond between atoms and molecules consisting of hydrogen

hydrogen bonds

an electrostatic force of attraction, “flirtation,” between a partial positive change that allows each hydrogen to be attracted to a nearby atom with a partially negative charge (oxygen in a water molecule for example)

• called hydrogen bonds because one atom in this attraction is always an hydrogen bond

  *usually represented by dotted lines

polar molecule

one or more sides are slightly positively charged and the other is slightly negatively charged

• unequal distribution of charges

  • polar molecules adjust or turn when in a magnetic field so that the side of the molecule with the positive charges are facing or attracted towards the plate with negative charges as it experiences an electric force in the direction of the field (and vice versa)

nonpolar molecule

the charges are distributed evenly throughout the molecule

• no “poles”

  • nonpolar molecules do not experience a force when in a magnetic field since their charges are evenly distribute throughout

intermolecular forces

attractions between molecules

• what holds molecules together

dipole-induced dipole

an intermolecular force that occurs between polar molecules and non polar molecules

• acts between a polar molecule and a nonpolar molecule as the dipoles of the polar molecule “induce” a temporary dipole in a nonpolar molecule which forms an attraction

london dispersion

an intermolecular force that is inevitable in attractions between any molecule, both polar and nonpolar (however, is the only intermolecular force between nonpolar molecules)

• acts between nonpolar molecules as they form temporary dipoles and become temporarily polar to form an attraction

  *because london dispersion does not rely on polarity, its strength is due to the size and shape of the molecules

chemical reactions

breaking chemical bonds and forming new ones

• matter cannot be created or destroyed, they can only be rearranged

  *our cells are constantly rearranging molecules

reactants

conversion of starting molecules in a chemical reaction

product

the material resulting from the chemical reaction

cohesion

the tendency of molecules of the same kind to stick together

• stronger for water than most other liquids due to hydrogen bonds between water molecules (though they only last for a few trillionths of a second)

  • ex. paper clip floating on top of the water versus sinking in

    • the water bonds together through hydrogen bonding which creates a film that holds up the paper clip

  • highly important in the living world

    • trees depend on cohesion to help transport water and nutrients from the roots to the leaves

surface tension

a measure of how difficult it is to stretch or break the surface of a liquid (for water, due to its hydrogen bonds)

• an extension of the concept on cohesion

adhesion

the clinging of one substance to another

• stronger for water than most other liquids due to hydrogen bonds between water molecules and other polar substances (though they only last for a few trillionths of a second)

  • ex. water travelling more up a capillary tube than alcohol despite gravity

    • water molecules are more polar than alcohol molecules and they form hydrogen bonds with other polar molecules (the glass in this case), resulting in a stronger attraction and climbing of water molecules up the tube

  • in plants, the thinness of a plant’s veins enhances the adhesion of water to its cell walls (and counters the force of gravity)

thermal energy

the energy associated with the random movement of atoms and molecules

heat

thermal energy in transfer from a warmer to a colder body of matter

temperature

a measure of intensity of heat

• the average speed of molecules in a body of matter

  • because of hydrogen bonding, water has a stronger resistance to temperature change than most other substances

high specific heat capacity

resistance to temperature change (due to hydrogen bonding for water)

• heat must be absorbed to break hydrogen bonds

  • to raise the temperature of water, hydrogen bonds between molecules must be broken before the molecules can move faster

    • water absorbs a large amount of heat, most of it disrupting hydrogen bonds, while warming up only a little

  • ex. warming up water and oil

    • the rate of temperature increasing for water is SLOWER than that of oil which is faster than water

    • water has a high heat capacity compared to oil, which means that water can take in more heat before the temperature of the substance increases

high specific heat of evaporation

when a substance evaporates (changing its physical state from a liquid to gas) and the surface of the liquid that remains cools down

• occurs because the molecules with the greatest energy (the hottest ones) leave

• heat is released when hydrogen bonds forms

  • when water cools, water molecules slow down and more hydrogen bonds form, releasing a considerable amount of heat

• ex. alcohol evaporating faster than water

  • water has a higher specific heat of evaporation as it requires more heat to increase the thermal energy to evaporate than alcohol

density

ice floats because it is less dense than water

• as water freezes, each molecule forms stable hydrogen bonds with other water molecules in a certain formation and creating a -D crystal

  • formation forces the atoms to be more spaciously apart in ice whereas in water, molecules are more tightly packed (concentration of mass in volume)

  • ex. ice floats on oil in which forms a layer between ice and water, and when ice melts, the water droplets sinks to the bottom

    • ice is less dense than water because of the arrangement of molecules with hydrogen bonds allow for it to be less packed with mass per volume than water which is more packed since hydrogen bonds are constantly breaking and reforming

  • if ice were to be more dense: ponds, lakes, and oceans would freeze

    *when water cools, ice serves as an insulating blanket that traps heat in water and prevents it from freezing (keeps marine life alive and serves as a ground for polar bears)

solution

a liquid consisting of an uniform mixture of two or more substances

solvant

the dissolving agent

aqueous

when water is the solvant

solute

a substance that is being or is dissolved

water as a solvant

water is the solvant of life which its versatility as a solvant results from the polarity of its molecules

• can dissolve both ionic compounds and polar molecules

ionic compounds in water

• when ionic compounds dissolve in water

  • the positively charged hydrogen ends of the water molecules form attractions to the negatively charged ends of the ionic compound

  • the negatively charged oxygen ends of the water molecules cling to the positively charged ends of the ionic compound

  • these processes work inward from the surface of the ionic compound in which water molecules eventually surrounded and separate all ions

    • ex. table salt — NaCl

polar molecules in water

• when polar molecules dissolve in water:

  • the positively charged hydrogen ends of water molecules are attracted to the negatively charged ends of the polar molecules

  • the negatively charged oxygen ends of water molecules are attracted to the positively charged ends of the polar molecules

  • these processes continue until all water molecules surround the compound and form hydrogen bonds with its polar regions

    • ex. sugar

dissociation of water molecules

water molecules breaking into hydrogen ions (H+) and hydroxide ions (OH-)

• in liquid water, a very small percentage of water molecules break apart into ions

  • however, these ions are highly reactive

  • changes in concentration can drastically affect a cell’s proteins and other complex molecules

    • some chemical compounds help add or remove H+ from an aqueous solution

acid

a substance that donates hydrogen ions to solutions

• has a higher concentration of H+ than OH-

  • ex. hydrochloric acid (HCl) is an acid in the gastric juice of the stomach

base

a substance that reduces the hydrogen ion concentration in a solution

• some bases like sodium hydroxide (NaOH) donates OH, which forms with H+ to form water molecules, thus reducing the concentration

  • this is a common ingredient in oven cleaners

  • other bases accept H+ from a solution, increasing -OH concentration

pH scale

describes how acidic or basic a solution is (pH means potential for hydrogen)

• ph 0 is the most acidic to pH 14 which is the most basic

  • decrease in pH equals increase in acidity (more H+ concentration)

  • increase in pH equals increase in basicity (more OH- concentration)

• each pH unit represents a 10-fold change in the H+ concentration of a solution

  • lemon juice at pH 2 has 10x more H+ than cola at pH3 and has 100x more H+ than tomato juice at pH4 (pH 7 is 10^-7 mol H+)

neutral pH

pH of 7

• pure water and aqueous solutions that are not acidic not basic are neutral

  • concentration of H+ and OH- are equal

  *the pH of most cells is close to 7 and blood pH is close to 7.4

  • a person cannot survive for more than a few minutes if blood pH drops to 7.0 or rises to 7.8

buffers

substances that minimize changes in pH

• in biological fluids to maintain homeostasis

  • accepts H+ when in excess or donates when depleted

carbon dioxide

main product of fossil fuel combustion

• the increasing the release of CO2 into the atmosphere is what causes climate change

  • 25% is absorbed by oceans (which seems like a good thing at first) but levels are rising and the increased absorption is harming marine life and ecosystems

ocean acidification

CO2 dissolving in seawaters lowers the pH of ocean water

• impact of lower pH on coral reefs

  • as seawater acidifies, extra H+ combine with carbonate ions (CO32-) to form bicarbonate ions (HCO3-)

    *this reduces the carbonate ion concentration available to corals and other shell-building animals for calcification (lower the concentration of carbonate ions, lower the rate of calcification, thus slows the growth of corals and animals)

calcification

coral animals combine calcium and carbonate ions to form their calcium carbonate skeletons

scientist research on ocean acidification

• controlled experiment (artificial habitat)

  • constant variables are the pH, temperature, calcium ions concentration

  • varied variable is the carbonate ion concentration in sea water

    *provided evidence that ocean acidification and its resulting reduction in carbonate ion concentration negatively affects coral reefs

• observation or study (natural habitat)

  • carbon dioxide is being released by underground volcanos which lowers the pH of seawater

    *reduction in coral diversity and shifts to a lesser structurally complex and slower growing coral reefs

• scientists synthesize conclusions using multiple lines of evidence

  • ocean acidification results from both where the pH naturally varies which has an impact on the health of coral reefs and the diversity of organisms they suppitt

    • a decrease in structural integrity equals a decrease in organisms that they can support (coral reeds are havens of organism diversity)

presence of water in the search of extraterrestrial life

the presence of water is important in the search for extraterrestrial life

• emergent properties of water supports life on earth (idea if life was able to form or exist in environments with water)

  • water plays important roles from moderating temperature to functioning as the solvant of life

• ex. mars

  • has ice caps at both pole and has signs that water may exist elsewhere on the planet