Biology Chapter 2 Part 1 - Tagged PDF

BIOLOGY 1 COMPOSITION OF MATTER Atoms Building blocks of elements Atoms of elements di5er from one another Subatomic Particles Nucleus Protons (p+) are positively charged Neutrons (n0) are uncharged or neutral Orbiting the nucleus Electrons (e–) are negatively charged Atoms are...

BIOLOGY 1 COMPOSITION OF MATTER Atoms Building blocks of elements Atoms of elements di5er from one another Subatomic Particles Nucleus Protons (p+) are positively charged Neutrons (n0) are uncharged or neutral Orbiting the nucleus Electrons (e–) are negatively charged Atoms are electrically neutral: Number of protons equals numbers of electrons in an atom Ions are atoms that have lost or gained electrons5 Subatomic Particles 6 (a) Hydrogen (H) (b) Helium (He) (c) Lithium (Li) (1p+; 0n0; 1e–) (2p+; 2n0; 2e–) (3p+; 4n0; 3e–) KEY: Proton Neutron Electron F I G U R E 2. 2 ATO M I C S T R U C T U R E O F T H E T H R E E S M A L L E S T ATO M S. 7 MOLECULES AND COMPOUNDS Molecule—two or more atoms of the same elements combined chemically Example: H (atom)  H (atom) H2 (molecule) (reactants) (product) Compound—two or more atoms of dierent elements combined chemically to form a molecule of a compound Example: 4H  C CH4 (methane) 8 CHEMICAL BONDS AND CHEMICAL REACTIONS - Electrons are distributed around the dense core of atoms in shells (orbitals): - 1st shell takes only 2 electrons to Cll up (stabilize). That is why atoms of helium He2 are stable (nonreactive or inert). - 2nd, 3rd, …7th shell: each takes 8 electrons to be complete. - Number of electrons in the last shell (valence shell) is key in determining the chemical bonding behavior of atoms 9 CHEMICAL BONDS AND CHEMICAL REACTIONS If valence shell of an atom has 2 (helium) or 8 (Neon, Argon, Xenon, …etc) electrons, the atom is stable (non-reactive or chemically inert) When the valence shell is less than 8, the atom tends to gain, lose, or share electrons with other atoms to complete their outermost orbital. Chemical bonds form. This will lead to: Atoms reach a stable state Bond formation produces a stable valence shell 10  (a) Chemically inert elements Outermost energy level (valence shell) complete 8e 2e 2e He Ne Helium (He) Neon (Ne) (2p+; 2n0; 2e–) (10p+; 10n0; 10e–) F I G U R E 2. 5 A C H E M I C A L LY I N E R T A N D R E A C T I V E E L E M E N T S. 11  (b) Chemically reactive elements Outermost energy level (valence shell) incomplete 4e 1e 2e H C Hydrogen (H) Carbon (C) (1p+; 0n0; 1e–) (6p+; 6n0; 6e–) 1e 6e 8e 2e 2e O Na Oxygen (O) (8p+; 8n0; 8e–) Sodium (Na) (11p+; 12n0; 11e–) F I G U R E 2. 5 B C H E M I C A L LY I N E R T A N D R E A C T I V E E L E M E N T S. 12 CHEMICAL BONDS Ionic bonds Form when electrons are completely transferred from one atom to another Allow atoms to achieve stability through the transfer of electrons Ions Result from the loss or gain of electrons Anions have negative charge due to gain of electron(s) Cations have positive charge due to loss of electron(s) Tend to stay close together because opposite charges attract 13  Chemical Bonds Ionic bonds + – Na Cl Na Cl Sodium atom (Na) Chlorine atom (Cl) Sodium ion (Na+) Chlorine ion (Cl–) (11p+; 12n0; 11e–) (17p+; 18n0; 17e–) Sodium chloride (NaCl) F I G U R E 2. 6 F O R M AT I O N O F A N I O N I C B O N D. 14 CHEMICAL BONDS Covalent bonds Atoms become stable through shared electrons Electrons are shared in pairs Single covalent bonds share one pair of electrons Double covalent bonds share two pairs of electrons Examples of atoms that can make covalent bond: C and C, C and H, O and O, O and H, N and H, etc. 15 Chemical Bonds Covalent bonds 16  Chemical Bonds Covalent bonds Reacting atoms Resulting molecules H H H C H C H or H H H Hydrogen atoms Carbon atom Molecule of methane gas (CH4) (c) Formation of four single covalent bonds F I G U R E 2. 7 C F O R M AT I O N O F C O VA L E N T B O N D S. 17 CHEMICAL BONDS COVALENT B ONDS Covalent bonds are either nonpolar or polar Nonpolar Electrons are shared equally between the atoms of the molecule (e.g.: C-C and C-H) Electrically neutral as a molecule Example: carbon dioxide (a) Carbon dioxide (CO2) 18 CHEMICAL BONDS COVALENT B ONDS Polar Electrons are not shared equally between the atoms of the molecule (one of the atoms have higher aHnity for the electrons than the other Molecule has a positive and negative side, or pole δ– Example: water δ+ δ+ (b) Water (H2O) 19 CHEMICAL BONDS Hydrogen bonds Form between atoms in molecules with polar covalent bonds Weak chemical bonds Hydrogen is attracted to the negative portion of a polar molecule Provides attraction between molecules Responsible for the surface tension of water Important for forming intramolecular bonds, as in protein structure 20  Hydrogen Bonds δ+ H H O δ– Hydrogen bonds δ+ δ+ δ– δ– δ– H H O O δ+ δ+ H H H δ+ O H δ– (a) (b) 21 F I G U R E 2. 9 H Y D R O G E N B O N D I N G B E T W E E N P O L A R WAT E R M O L E C U L E S. PATTERNS OF CHEMICAL REACTIONS Synthesis reaction (A  B AB)  Atoms or molecules combine  Energy is absorbed for bond formation  Underlies all anabolic activities in the body Decomposition reaction (AB A  B)  Molecule is broken down  Chemical energy is released  Underlies all catabolic activities in the body Exchange reaction (AB  C AC  B and AB  CD AD  CB)  Involves both synthesis and decomposition reactions as bonds are both made and broken In general, most chemical reactions are reversible 22  CHEMICAL REACTIONS Absorb energy Release energy Absorb & release energy Anabolic Catabolic Anabolic / catabolic (condensation) 23 Factors inIuencing rate of chemical reactions 24 TA B L E 2. 4 FA C T O R S I N C R E A S I N G T H E R AT E O F C H E M I C A L R E A C T I O N S. BIOCHEMISTRY: CHEMICAL COMPOSITION OF LIVING MATTER Chemicals found in the body are either Inorganic or Organic compounds -Inorganic compounds: mostly lack carbon and/or small; e.g., CO2, H2O, NH3, NaCl, Calcium, etc. Inorganic compounds in biological systems include water, salts, acids, and bases - Organic compounds: contain Carbon, tend to be larger in size than inorganic compounds; e.g., glucose (C6H12O6), lipids, proteins, etc. Organic compounds in living matter include proteins, lipids, sugars and nucleic acids25  INORGANIC COMPOUNDS Water : All properties are related to its ability to form hydrogen bonds High Heat capacity: amount of heat required to raise the temperature of one mole or one gram of a substance by one degree Celsius without change of phase. This means that water can absorb or release large amount of heat before changing temperature suddenly (sun exposure, winter, etc.) Polarity/solvent properties: because of its polarity, water is the most common solvent for inCnite number of solutes; it is a Universal solvent. 26  INORGANIC COMPOUNDS Water properties (cont.) Capillary action: ability of water to Iow in narrow spaces without the assistance of, and in opposition to external forces like gravity High Surface tension: increased cohesion of water molecules to each others. Think of the ability of some insects (e.g., water striders) to run on the water surface Chemical Reactivity: Water is an important reactant in many types of chemical reactions. Hydrolysis is the breakdown of large molecules due to addition of water into the reaction Cushioning /protection: Iuids (water solutions) around heart, brain (CSF), developing fetus, 27 INORGANIC COMPOUNDS Salts: Ionic compounds that contain cations (other than H+) and anions (other than OH-) Salts of many metals are present in our bodies. e.g.: calcium salts (found in bone & teeth) and phosphorous (found in bone, teeth, nucleic acid) Easily dissolve (dissociate) in solvents like water Salts dissolved in body Iuids to form ions (electrolytes= conductors of electricity). e.g.: Sodium Chloride (Na+ and Cl-) is an example of electrolyte in blood Electrolyte (ionic) balance is very important for the normal functioning of the body  loss of 28 INORGANIC COMPOUNDS Acids: Substances that release H+ (protons) Taste sour, very corrosive and damaging. Acids that ionize completely and release their protons are called strong acids (HCl) HCl  H+ + Cl- (complete dissociation) (proton) (anion) Acids that dissolve partially are called weak acids (e.g., carbonic acid). H2CO3  H+ + HCO3- + H2CO3 (partial dissociation) (proton) (anion) 29  100 HCl 100 H+ + 100 Cl- 100 NaOH 100 Na+ + 100 OH- H+ 100 H Cl 100 100 Cl- pH= - log [H+] = - log 30 INORGANIC COMPOUNDS Bases: Substances that accept H+ (protons)  release OH- to accept H+ and form water Taste bitter and have a slippery feeling to them, NaOH  Na+ + OH- cation hydroxyl ion Bases that dissociate completely and release all their OH- are considered strong bases (NaOH) HCO3- can function as a weak base as it accepts H+ Neutralization Reaction: Type of exchange 31 INORGANIC COMPOUNDS The pH concept - The concentration of protons [H+] in solutions (in blood, body Iuids, etc.) reIects the degree of acidity or alkalinity of the solution. - It is not easy to directly measure proton [H+] concentration in a solution. -But as protons are ions (conductors of electricity), we can indirectly measure their relative concentration by measuring their degree of conductivity in the solution. -The output (reading) is called the pH = which is a measure of proton activity ([H+]) in a 32 INORGANIC COMPOUNDS The pH concept (cont.) -Example 1: if [H+]= 0.00001=10-5, then pH= -Log [10-5]  -(-5)=5 -Quiz 1: if [H+]= 0.000001, then pH=? -Example 2: If pH=3, then–log H+=3[H+]=10-3= 0.001 -Quiz 2: if pH=8, then H+ = ? A di5erence of 1 unit = 10x up- or down- change in [H+]; A change in pH from 4 to 5 = 10-fold decrease in [H+] A change in pH from 6 to 5 = 10-fold increase in [H+] 33 INORGANIC COMPOUNDS The pH concept (cont.) pOH is also a measure of proton activity in a solution. Example: if [OH-] = 10-7, then pOH = -Log [OH-]  - (-7) = 7 Hence, pH = 14 – 7 = 7 (neutral solution) Range of pH scale = Zero (very acidic) to 14 (very basic) pH 7 is basic or alkaline 34 The pH scale and pH values of representative substances 35 INORGANIC COMPOUNDS Bu5ers In biological systems, range of [H+] (or proton activity) is narrow. Increased or decreased [H+] can disrupt or damage biological molecules, halt metabolism or kill cells. Example: [H+] in blood is around = 10-7 M; any sudden change in blood [H+] could be fatal  needs to be regulated Bu5ering systems exist to prevent sudden or dramatic changes in [H+] Weak acids and weak bases are good bu5ers 36