Topic 2: The Chemical Basis of Life PDF

Topic 2 The Chemical Basis of Life PowerPoint Lectures for Biology, Seventh Edition Neil Campbell and Jane Reece Lectures by Chris Romero Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Learning Objectives (LOBs) 1. Define molecules, compounds and elements and i...

Topic 2 The Chemical Basis of Life PowerPoint Lectures for Biology, Seventh Edition Neil Campbell and Jane Reece Lectures by Chris Romero Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Learning Objectives (LOBs) 1. Define molecules, compounds and elements and identify key chemical elements and their properties. 2. Define isotopes and radioisotopes and provide examples of their clinical applications. 3. Describe and compare ionic bonds, covalent bonds, hydrogen bonds and Van der Waals interactions. Reading: Campbell Biology Chapter 2 Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Elements and Compounds Organisms are composed of matter Matter: anything that takes up space and has mass Matter is made up of elements Elements: substances that cannot be broken down to other substances by chemical reactions 92 elements exist in nature (listed in the Periodic table) Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  The Periodic Table Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Compounds Compound: – a substance consisting of 2 or more elements combined in a fixed ratio – has characteristics different from those of its elements + Figure 2.2 Sodium (Na) Chlorine (Cl) Sodium Chloride (NaCl) Metal Copyright © 2005 Pearson Education, Inc. Poisonous gas publishing as Benjamin Cummings Table salt  Essential Elements of Life Essential elements: – Include carbon (C), hydrogen (H), oxygen (O), and nitrogen (N) – Make up 96% of living matter A few other elements make up most of the remaining 4% of living matter (Ca, K, Na, etc.) Trace elements (0.01% of living matter) are required by an organism in only minute quantities: e.g. Fe, I, Se, Cu Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  The effects of trace element deficiencies Trace elements → required by an organism in minute quantities but are very important! This woman suffers from iodine deficiency (goiter: enlargement of thyroid gland) Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Trace element deficiencies Iodine (I) deficiency: Hypothyroidism - Iodine → essential trace element for production of the thyroid hormones (thyroxine and triiodothyronine) Iron (Fe) deficiency: Iron deficiency anaemia - Fe is part of the haemoglobin structure Heme group Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Clinical correlations: copper deficiency (copper transporter disorder) Batzios S, Tal G, Di Stasio AT, Peng Y, Charalambous C, Nicolaides P, Kamsteeg EJ, Korman SH, Mandel H, Steinbach PJ, Yi L, Fair SR, Hester ME, Drousiotou A, Kaler SG. Newly identified disorder of copper metabolism caused by variants in CTR1, a high- affinity copper transporter. Human Molecular Genetics. 2022; 00 (00): 1–10. An element’s properties depend on the structure of its atoms Each element consists of a specific atom Atom: – The smallest unit of matter that still retains the properties of an element – An element and its atom have the same symbol Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Subatomic Particles Atoms of each element are composed of smaller parts called subatomic particles Subatomic particles include: – Neutrons: no electrical charge (0) – Protons: positively charged (+1) – Electrons: negatively charged (-1) Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  The mass of subatomic particles Protons and neutrons are identical in mass – Their mass is 1.7x10-24 g = 1 Dalton The mass of electrons is considered negligible – Electrons have mass which is 1/2000 that of protons and neutrons Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Simplified models of an atom Protons and neutrons are found in the atomic nucleus Electrons surround the nucleus in a “cloud” Cloud of negative Electrons charge (2 electrons) Nucleus (a) This model represents the (b) In this even more simplified electrons as a cloud of model, the electrons are negative charge, as if we had shown as two small blue taken many snapshots of the 2 spheres on a circle around the electrons over time, with each nucleus. dot representing an electron‘s Figure 2.4 position at one point in time. Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  The atom: structure Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Atomic Number Atoms of the various elements differ in their number of subatomic particles The atomic number (Z) of an element is: – the number of protons (Z=p) – equal to the number of electrons (Z=e) – unique to each element – written as a subscript to the left of the symbol Mass number A Atomic number X Z Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Atomic mass = Mass number The mass number (A) of an element: – the sum of protons plus neutrons in the nucleus of an atom (Α = p + n) – an approximation of the atomic mass of an atom (in Dalton) – written as a superscript to the left of the symbol Mass number 4 He 2 Atomic number Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Atomic and Mass number ▪ Αtomic number (Ζ) = the number of protons (p) of an atom Z=p ◼ Μass number (Α) = the sum of neutrons (n) and protons (p) present in the nucleus of an atom Α=p+n → number of neutrons = mass number (protons + neutrons) - atomic number (protons) Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Atomic and Mass number Mass number =number of protons + 23 number of neutrons 11 Na Atomic number =number of protons Number of protons: 11 Number of electrons: 11 Number of neutrons: 23-11=12 Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Atomic mass = Mass number Atomic mass: the atom’s total mass, equal to the mass number (in Daltons) Mass number = number of protons + 23 number of neutrons 11 Na Atomic Mass = Mass number in Daltons= 23 Daltons Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Isotopes Isotopes: Different forms in which atoms of a given element may occur in – Differ in the number of neutrons in the atomic nucleus (have different mass number) – Have the same number of protons (same atomic number) Radioactive isotopes (radioisotopes): – Unstable → spontaneously decay to give off particles and energy – Used in nuclear medicine, imaging etc → e. g C14, H3 Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Isotopes examples Hydrogen isotopes: http://www.youtube.com/watch?v=Jdtt3LsodAQ&feature=related Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Clinical applications of radioisotopes Disease diagnosis (imaging) Disease treatment (e.g I131 for thyroid cancer, hyperthyroidism) Assessment of degree of disease severity Treatment monitoring Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Nuclear medicine imaging Example: PET scan - Positron Emission Tomography - Based on the use of radioactive isotopes - Imaging technique that produces a 3D image of functional processes in the body - Detection of γ-rays emitted by a radioisotope (tracer) introduced into the body as part of a biologically active molecule (e.g. F18-FDG fluoro-deoxy-glucose) Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Radioactive Isotopes Imaging Cancerous throat tissue Figure 2.6 Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  PET scan Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Electron Configuration and Chemical Properties The chemical behavior of an atom is defined by its electron configuration Electron configuration: the distribution of an atom’s electrons in shells The periodic table of the elements shows the electron distribution for all the elements Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  The Periodic Table Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  The formation of chemical bonds between atoms Valence electrons: – The electrons in the outermost shell, called valence shell – Determine the chemical behavior of an atom – Involved in chemical bonds between atoms Atoms either share or transfer valence electrons This results in formation of chemical bonds between atoms Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Electron shells Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Periodic table Hydrogen 2 Atomic number Helium He 2He 1H 4.00 Element symbol Atomic mass First Electron-shell shell diagram Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon 3Li 4Be 3B 6C 7N 8O 9F 10Ne Second shell Sodium Magnesium Aluminum Silicon Phosphorus Sulfur Chlorine Argon 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar Third shell Figure 2.8 Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Chemical bonds Covalent bonds: sharing of electrons between atoms Ionic bonds: transfer of electrons from one atom to another Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Covalent Bonds A covalent bond is the sharing of a pair of valence electrons by two atoms Hydrogen atoms (2H) 1 In each hydrogen atom, the single electron is held in its orbital by + + its attraction to the proton in the nucleus. 2 When two hydrogen atoms approach each other, the electron of + + each atom is also attracted to the proton in the other nucleus. 3 The two electrons become shared in a covalent bond, + + forming an H2 Hydrogen molecule. molecule (H2) Figure 2.10 Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Molecules and covalent bonds A molecule consists of 2 or more atoms held together by covalent bonds A single bond: the sharing of one pair of valence electrons A double bond: the sharing of two pairs of valence electrons Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Single and double covalent bonds Name Electron- Structural Space- (molecular shell formula filling formula) diagram model (a) Hydrogen (H2). Two hydrogen H H atoms can form a single bond. (b) Oxygen (O2). Two oxygen atoms share two pairs of O O electrons to form a double bond. Figure 2.11 A, B Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Covalent bonding in compounds Name Electron- Structural Space- (molecular shell formula filling formula) diagram model (c) Water (H2O). Two hydrogen atoms and one O H oxygen atom are joined by covalent H bonds to produce a molecule of water. (d) Methane (CH4). Four hydrogen atoms can satisfy H the valence of one carbon H C H atom, forming methane. H Figure 2.11 C, D Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Covalent bonds Electronegativity: – the tendency of an atom to attract electrons towards itself – is affected by both its atomic number and the distance of its valence electrons from the charged nucleus (electron shells) The more electronegative an atom → the more strongly it pulls shared electrons toward itself http://www.youtube.com/watch?v=y2uUl7fyPBM&feature=related Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Electronegativity Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Electronegativity Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Covalent bonds In a non-polar covalent bond: – The atoms have similar electronegativities – Share the electrons equally In a polar covalent bond: – The atoms have different electronegativities – Share the electrons unequally Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Covalent bonds Polar covalent bond example: H20 Because oxygen (O) is more electronegative than hydrogen (H), shared electrons are pulled more toward oxygen. – This results in a partial negative charge on the oxygen and a O partial positive charge on the hydrogens. Figure 2.12 H H + + H2O Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Ionic Bonds In ionic bonds, electrons are transferred from one atom to another Transfer of electrons between two atoms creates ions Ions: – charged atoms – atoms with more or fewer electrons than usual Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Ions: anions and cations Types of ions: - Anions: negatively charged ions - Cations: positively charged ions Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Ionic bond An ionic bond is an attraction between anions and cations 1 The lone valence electron of a sodium 2 Each resulting ion has a completed atom is transferred to join the 7 valence valence shell. An ionic bond can form electrons of a chlorine atom. between the oppositely charged ions. + – Na Cl Na Cl Na+ Cl– Na Cl Sodium ion Chloride ion Sodium atom Chlorine atom (a cation) (an anion) (an uncharged (an uncharged Figure 2.13 atom) atom) Sodium chloride (NaCl) Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Ionic compounds Ionic compounds are often called salts, which may form crystals e.g. NaCl Na+ Cl– Figure 2.14 http://www.youtube.com/watch?v=xTx_DWboEVs Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings Ionic vs covalent bond http://www.youtube.com/watch?v=QqjcCvzWwww&feature=related Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings Strong vs weak chemical bonds Strong Chemical Weak Chemical Bonds Bonds Covalent Bonds Hydrogen Bonds Ionic Bonds Van der Waals interactions Bonds between Bonds between atoms atoms or molecules Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings Weak Chemical Bonds Several types of weak chemical bonds (interactions) are important in living systems Examples: - Hydrogen bonds - Van der Waals interactions Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Hydrogen bonds Interaction between neighbouring molecules Form between a hydrogen atom (Η) covalently bonded to an electronegative atom of a molecule and an electronegative atom (oxygen, nitrogen or fluoride) of another molecule e.g. hydrogen bonds between neighbouring water molecules (Η2Ο) →The H atom of water molecule forms hydrogen bond with the O atom of another water molecule Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Hydrogen Bonds e.g. The H atom of water molecule forms hydrogen bond with the N atom of ammonia – + A hydrogen Water H bond results (H2O) O from the attraction between the H partial positive + charge on the hydrogen atom – of water and the partial Ammonia negative charge N (NH3) on the nitrogen H H atom of + H + ammonia. Figure 2.15 + Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings Ionic and hydrogen bonds in NaCl solution Ionic bonds formed between Na+ and Cl- Η2Ο (water molecule) Hydrogen bonds formed between water molecules Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Van der Waals Interactions Attractive forces → developed between molecules or atoms of a molecule (other than covalent bonds, ionic bonds and hydrogen bonds) Occur when transiently positive and negative regions of molecules attract each other Weaker than H-bonds e.g. found in polypeptide tertiary structure → protein folding van der CH Waals CH2 H3 CHCH O2 interactions Hydrogen H3C C3H H C bond O CH 3 Polypeptide HO C backbone CH CH SS CH 2 2 2 O CHNH- CH + OC 2 3 2 Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Molecular Shape and Function Role of weak chemical bonds: – Reinforce the shapes of large molecules – Help neighbouring molecules adhere to each other The precise shape of a molecule is important to its function in the living cell (e.g. protein shape) Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Molecules vs Compounds Covalent bonds can form between atoms of the same element or atoms of different elements A molecule: consists of 2 or more atoms held together by covalent bonds e.g. H2 A compound: is a combination of 2 or more different elements e.g. H20 Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings Chemical reactions Chemical reactions are the making and breaking of chemical bonds The starting molecules of a chemical reaction are called reactants The final molecules of a chemical reaction are called products Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Chemical reactions Convert reactants to products + 2 H2 + O2 2 H2O Reactants Reaction Product Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings  Chemical reactions Photosynthesis is an example of a chemical reaction: CO2 + H2O light O2 + C6H12O6 Figure 2.18 Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings Summary Molecules, compounds and elements, trace element deficiencies Isotopes and radioisotopes: role of radioisotopes in imaging Chemical bonds: ionic bonds vs covalent bonds, hydrogen bonds and Van der Waals interactions Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings SBA (Single Best Answer) example A patient is recently diagnosed with anaemia caused by a trace element deficiency. Which is the trace element responsible for the anaemia? A. Iodine B. Selenium C. Iron D. Zinc E. Manganese Copyright © 2005 Pearson Education, Inc. publishing as Benjamin Cummings