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CHAPTER 2

CHAPTER 2. CHEMISTRY AND BIOCHEMISTRY. TOPIC 2: BASIC CHEMISTRY FOR BIOLOGY. Matter consists of chemical elements in pure form and in combinations called compounds An element ’ s properties depend on the structure of its atoms

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CHAPTER 2

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  1. CHAPTER 2 CHEMISTRY AND BIOCHEMISTRY

  2. TOPIC 2: BASIC CHEMISTRY FOR BIOLOGY • Matter consists of chemical elements in pure form and in combinations called compounds • An element’s properties depend on the structure of its atoms • The formation and function of molecules depend on chemical bonding between atoms • Chemical reactions make and break chemical bonds

  3. Chemistry is the science dealing with the properties & the transformations (chemical reactions) of all forms of matter Matter is ….anything that takes up space (volume) or has mass. Matter is made of elements.

  4. Concept 1 Matter consists of chemical elements in pure form and in combinations called compounds

  5. Matter: Anything that takes up space and has mass States of Matter

  6. ELEMENT = A substance that cannot be broken down into other substances by chemical 
    reactions. 
   • TRACE ELEMENT= Element required by an organism in extremely minute quantities. Example: Fe (Iron) • COMPOUND= A pure substance composed of two or more elements combined in a fixed 
 ratio. 
Example: NaCI (sodium chloride)

  7. All Matter is composed of : Atom – Means “unable to be cut” (Greek: atomos) Smallest possible unit of matter that retains the physical and chemical properties of 
its element.

  8. Elements • Substances that can’t be broken down into simpler substances • 92 naturally occurring elements • 25 - 36 chemical elements are essential to life. • 96% of human mass is made up of C, H, O, N

  9. C H N O P S Structure of Atoms

  10. Structure of Atoms Chemistry is the science dealing with the properties & the transformations (chemical reactions) of all forms of matter Matter is ….anything that takes up space or has mass. Matter is made of elements.

  11. Structure of Atoms Atoms contain protons, neutrons, and electrons. Protons - positive charge – 1 atomic mass unit Neutrons – no charge – 1 atomic mass unit Electrons – negative charge – almost no mass

  12. Structure of Atoms

  13. Structure of Atoms Atomic Number = number of protons In atom Atomic Mass = mass of protons, neutrons and electrons in atom

  14. Structure of Atoms C H N O P S

  15. Isotopes

  16. Structure of Atoms Isotopes – alternative forms of an atom that differ in their number of neutrons (same atomic number, different MASS number) How are these atoms different from each other? What can isotopes be used for?

  17. Isotopes ISOTOPES Atoms of an element that have the same atomic number but different 
 mass number. 
       They have the same number of protons, but a different number of neutrons. RADIOACTIVE ISOTOPEUnstable isotope in which the nucleus spontaneously decays, emitting 
 subatomic particles and/or energy as radioactivity. Loss of nuclear particles may transform one element to another 
 A radioactive isotope has a fixed half-life: HALF LIFE = Time for 50% of radioactive atoms in a sample to decay.

  18. Isotopes Differing Numbers of Neutrons 1 Hydrogen: H 1 1 Deuterium: H 2 1 H Tritium 3

  19. Applications of isotopes: Carbon Dating

  20. Applications of isotopes: Carbon Dating • 99% of carbon atoms have 6 neutrons (12C). • Most of the remaining 1% of carbon atoms have 7 neutrons (13C) while the rarest carbon isotope, with 8 neutrons, is 14C. • Both 12C and 13C are stable isotopes while 14C is radioactive • When 14C decays, one of its neutrons is converted to a proton and an electron. • This converts 14C to 14N, transforming the atom to a different element.

  21. Applications of isotopes: Carbon Dating By comparing the ratio of radioactive (14C) and stable (12C) isotopes in a fossil 
with the ratio of isotopes in living organisms, one can estimate the age 
of a fossil less than 50,000 years old.

  22. Application of isotopes: Nuclear scintigraphy • Diagnostic radiopharmaceuticals – Technetium 99 • Iodine 131: diagnosis and treatment of thyroid tumours • PET scanning (positron Emmision Tomography) • CT scanning (computer tomography) Generally short-lived isotopes linked to chemical compounds to be traced.

  23. Structure of Atoms: Electron Orbitals There are certain spaces around the nucleus where electrons are likely to be found. The electron is found a certain distance from the nucleus called a shell. Within a shell is an orbital that holds 2 electrons.

  24. Structure of Atoms: electron Orbitals The outermost electron shell is referred to as the valence shell.

  25. Structure of Atoms: Electron Orbitals RULES ABOUT ELECTRONS! 1. First energy shell can contain a maximum of 2 electrons 2. All other shells can contain a maximum of 8 electrons (there are exceptions) An atom is moststable when the valence shell contains a full 8 electrons = OCTET RULE

  26. Structure of Atoms RULES ABOUT ELECTRONS! Electrons determine chemical reactivity. Atoms are reactive unless they have a full valence shell of electrons!

  27. Chapter 2 Chemical Principles Structure of Atoms

  28. Chapter 2 Chemical Principles Structure of Atoms Which atoms on the periodic table have a full valence shell? Which atoms on the periodic table have only one electron in the valence shell? Which atoms on the periodic table have seven electrons in the valence shell?

  29. A Brief Review: Types of Bonding INTRAMOLECULAR FORCES • Ionic Bonding • Covalent Bonding • (Polar Covalent bonding) INTER-MOLECULAR FORCES • Dipole-dipole interactions (Van Der Waals Forces) • Hydrogen Bonds

  30. Bonding between elements (in compounds and molecules) • BONDING • IONIC BONDING COVALENT BONDING • non-polar covalent • and • polar covalent

  31. Bonding and electronegativity Electronegativity is the tendency of an atom to attract electrons. If the electronegativity of an atom is high, then it attracts and holds on to electrons. If the electronegativity of an atom is low, then it tends to give electrons away.

  32. Bonding and electronegativity Where on the Periodic Table do you find the atoms with high electronegativity? Where do you find those atoms with low electronegativity? Which atom has the highest electronegativity?

  33. Ionic Bonding Ionic bonding – Atoms transfer or accept electrons from one another Look at the following diagram with regards to sodium (Na) and chloride (Cl) What must each do to fulfill the octet rule?

  34. Ionic

  35. Bonding Ionic bonding The sodium ion and the chloride ion will be attracted to each other and form an ionic bond. Na + Cl - The ionic bond is due to the attractive forces between the now positively charged sodium & the negatively charged chloride. By looking at the periodic table, can you predict which atoms may form ionic bonds?

  36. Bonding: Covalent Bonding Non-polar covalent bond – equal sharing of electrons Polar covalent bond – unequal sharing of electrons What does the word polar mean?

  37. Covalent bonding

  38. Bonding: Covalent Bonding H C How many electrons are needed to fill the valence shell of carbon? How many electrons are needed to fill the valence shell of hydrogen? In the molecule of methane (CH4), the electrons are shared equally between the carbon and hydrogens.

  39. Bonding Covalent bonding

  40. Covalent Bonding

  41. Polar Covalent Bonding Polar covalent bond – unequal sharing of electrons A great example of a molecule with polar covalent bonds is water. Why is water considered polar? What is a partial positive and partial negative charge?

  42. Polar covalent bond

  43. Chapter 2 Chemical Principles Bonding Covalent bonding Polar covalent bond

  44. Inter-molecular forces: dipole-dipole interactions and hydrogen bonding

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