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Chapter 2 Notes

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by: Michelle Notetaker

Chapter 2 Notes Bio 1510

Michelle Notetaker
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Atoms, Properties of Water, Acids vs. Bases
Basic Life Mechanisms
Dr. Nataliya Turchyn
Class Notes




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This 6 page Class Notes was uploaded by Michelle Notetaker on Tuesday January 12, 2016. The Class Notes belongs to Bio 1510 at Wayne State University taught by Dr. Nataliya Turchyn in Summer 2015. Since its upload, it has received 26 views. For similar materials see Basic Life Mechanisms in Biology at Wayne State University.


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Date Created: 01/12/16
Chapter 2  The Nature of Molecules and Properties of Water  Atomic Structure o An atom is the smallest unit of an element o In an electrically neutral atom, the number of protons = to the number of electrons o Most atoms are composed of protons, electrons (e­), and neutrons.  Protons are positively charged  Neutrons have no charge  Electrons have negatively charge (e­) o Nucleus does NOT have electrons, but the electrons are moving around it. o Electrons are not flying away because electrons are negatively charged and are  attracted to the positively charge nucleus. o Hydrogen, oxygen, carbon, and nitrogen  Hydrogen is the smallest one  It only contains one proton  It has no neutrons  How are atoms categorized? o The symbol is the name of the element  Ex. C  Carbon, Na = Sodium, K = Potassium o Atomic number = the number of protons o Atomic mass = # of protons + # of neutrons o 1 amu (Dalton) is the mass in which atoms are measured  Proton = 1 amu  Neutron = 1 amu  Electron = roughly 0 amu o Atomic weight  depends on the force of gravity  Has same numeric number, but NOT the same units o Number of neutrons = atomic weight – atomic number  Isotopes o Atoms with the same number of protons and electrons, but a different number of  neutrons o Can be unstable  Nucleus breaks apart  Radioactive  o Carbon – 12  Atomic weight = 12  Protons = 6  Electrons = 6  Neutrons = 6 o Carbon – 13  Atomic weight = 13  Protons = 6  Electrons = 6  Neutrons = 7 o Carbon – 14 (unstable isotope)  Atomic weight = 14  Protons = 6  Electrons = 6  Neutrons  It’s nucleus has too many neutrons than protons  It breaks apart that releases lots of energy (radioactive decay) o Radioactive isotopes  when an isotope breaks apart it releases energy  Used for cancer or diagnosed disease  It can be used to predict the age of fossils up to 50,000 years old  “Carbon Dating” o The level of 14C in fossils is compared to the level of 14C in atmosphere o Carbon dating uses 14C to predict the age of fossils (remains of living organisms)  up to 50,000 years old o The level of 14C in the organisms is composed to the 14C in the atmosphere o All organisms take in 14C o When organism dies they stop taking in 14C and decays into 14N o The half­life is 5,730 years  It means that it takes 5,730 years for 50% of 14C to decay to 14N o The longer the organism is dead the less 14C, but more 14N o Not completely accurate but is a rough estimation o If the fossil is older than 50,000 years, then a different radioactive isotope must be used o Uncontrolled exposure can change genetic material or kill the cells  Atoms contain discrete energy levels o Of the 3 subatomic particles, only electrons are directly involved in chemical  activity o Electrons have different energy levels (shells) such as K, L, M, N…  K<L<M<N (K has the least amount of energy and N has the most)  Indicate the amount of energy that an electron has  Electrons farther from nucleus have more energy o The # of electrons in the outermost shell, called the valence shell, determines the  chemical properties of the atom o Atoms whose valence shells are not full have unpaired electrons and tend to  interact with other atoms, participating in chemical reactions o Only electrons determine the chemical properties of the atom o Unreactive if all electrons are paired already  Electrons occupy orbitals o Electrons orbit around the nucleus in orbitals such as s, p, d, f…  Areas around the nucleus where electrons are most likely to be found  No orbital can contain more than 2 electrons o Orbital  areas where most likely the electrons will be found  They come in various shapes  S is a sphere  P is a dumbbell  D is an X o Energy level K has a single s orbital (1s) and can hold up to two electrons o Energy level L has one s orbital (2s) and 3 p orbitals (2px, 2py, and 2pz), and can  hold up to eight e­ o Level K is a single 1s = 2 electrons o Energy level L is 2  shell it has a single 2s and 3 different orbitals  It has 4 different orbital = electrons o Level M = 18 electrons o 2n^2, n=the number of shell(s) o K = 1  shell o L = 2  shell rd o M = 3  shell o N = 4  shell  When electrons gain or lose energy… o Electrons can absorb light energy (photons) and move to a higher energy level o Electrons can lose or emit energy (heat or light) and move closer to the nucleus o Ion forms when an atom gains or loses the electron(s)  Cations = positively charged ions resulting from the loss of electron(s)  Anions = negatively charged ions resulting from the gain of electron(s) o Movement of electrons from one atom to another  redox reaction o OIL RIG = oxidation is losing/reduction is gaining o When atoms lose electrons, they lose energy, their electrons move closer to the  nucleus to lower energy level o When atoms gain electrons, they gain energy. Their electrons move away from  the nucleus which makes it a higher energy level  Chemical Bonds o Atoms with incomplete outer shells tend to react so that both atoms end up with  completed outer shells o These atoms may react with each other by sharing, donating, or receiving  electrons o These interactions usually result in atoms staying close together, held by  attractions called chemical bonds o Covalent bond, ionic bond, hydrogen bond, hydrophobic interaction, and van der  Waals attraction (strongest to weakest)  Ionic Bonds o Form between two oppositely charged ions (cations and anions)  Sodium  cation  It has more protons  Sodium undergoes oxidation which becomes a cation  Chlorine  anion  Chlorine gains the electron and undergoes reduction  Covalent Bonds o Two atoms share one or more pairs of outer­shell electrons o Can be single, double, and triple  Single covalent bond only shares 1 pair of electrons (or 2 electrons)  Double covalent bond shares 2 pairs of electrons (or 4 electrons)  Triple covalent bond shares 3 pairs of electrons (or 6 electrons) o Triple is the strongest, single is the weakest  Nonpolar Covalent Bonds o Electronegativity – atom’s attraction (pull) for shared electrons  o Nonpolar covalent bonds form between atoms that equally share their outer shell  electrons  Polar Covalent Bonds o Form between atoms that do NOT share their electrons equally because they have very different electronegativities  o Oxygen has a partial negative change o Hydrogen has a partial positive change  Hydrogen Bonds o Form between hydrogen atom of one molecule and an electronegative atom (O,N,  or F) of another molecule o Water is polar because it has 2 covalent bonds  H is FON  H=Hydrogen  F=Florine  O=Oxygen  N=Nitrogen  Chemical Reaction o The formation and breaking of chemical bonds o The rate of the chemical reaction depends on:  Catalysts   Are enzymes that speed the rate of reactions  Temperature  More chemical reactions proceed faster at higher temperatures  Concentrations of reactants vs. products  The more reactants, the faster the products will be created  Importance of Water o Water is cohesive  Cohesion  water molecules stick to other water molecules by hydrogen  bonding o Surface tension  Water strider walks on the surface of the pond without breaking the  water’s surface (surface tension) o Water is adhesive  Adhesion  water molecules stick to other polar molecules by hydrogen  bonding   Capillary action  Occurs when adhesion of water molecules to the glass is stronger  than cohesion of water molecules to each other  Both cohesive and adhesive properties of water work together to move  water up from root to shoots o Water has a high specific heat  A large amount of heat (energy) is needed to raise the temperature of  water o Water has a high heat of vaporization   A lot of heat (energy) is necessary to turn liquid water into vapor  Allows living things to release excess body heat via sweating o Frozen ice is less dense than liquid water  Bodies of water freeze from the top down  Ice floats on liquid water  Water molecules are more organized  Hydrogen bonds are more stable  Liquid water – hydrogen bonds are dynamic (less organized) o Water is a universal solvent  It dissolves many different substances  Solutes are substances that are being dissolved in water  Hydration shells = clouds of water molecules  Prevents sodium and chlorine to come together o Water organizes nonpolar molecules  Hydrophilic  “water loving”  Polar  Hydrophobic “water fearing”  Non­polar  Water causes hydrophobic molecules to aggregate or assume specific  shapes  Hydrophobic exclusion is when hydrophobic molecules such as oil  exclude themselves from water o Can form ions  Acid [H+]>[OH­]  Ex. Lemon juice and stomach  Base [OH­]>[H+]  Ex. Household bleach/small intestine  When water forms ions: Acids vs. Bases o Pure water = pH 7 o pH<7 acidic   The lower of the pH the stronger the acid o pH>7 basic  The higher the pH the stronger the base o Normal blood pH is 7.4  Why do we need buffers? o Buffer is a substance that minimizes changes in pH o Acts by donating H+ when the solution becomes too basic and accepting H+  when it becomes too acidic  Carbonic acid­bicarbonate buffer (blood)  Carbonic acid is a weak acid   Acidosis vs. Alkalosis o Acidosis is when blood becomes acidic due to higher concentration of H+ ions  Hypoventilation = when we breathe too little (pneumonia or emphysema)  Must decrease [H+]  Add bicarbonate ions (HCO3­)  Reduce [CO2] by inhaling more o Alkalosis  when blood pH becomes to basic due to lower concentration of [H+]  Hyperventilation = when we breathe too much (stress and anxiety)  Increase the concentration of H+  Use paper bag and start breathing  Increasing the concentration of CO2 in the blood  Add H2CO3 (carbonic acid) dissociates to H+ and returns blood back to  normal


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