1.1 Atoms And Bonding
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1.1 Atoms And Bonding

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An introduction and review of concepts required for success in year 12 biology (BIO4U).

An introduction and review of concepts required for success in year 12 biology (BIO4U).

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  • 1. Metabolic Processes
    • Unit 1
  • 2. Metabolic Processes
    • All organisms require energy to survive
    • Cells in an eukaryotic organism contain organelles (i.e. mitochondria), that transform the energy in food into energy that can be used for various cellular processes
    • Metabolic processes involve all the chemical reactions that take place in cells, as well as the chemical reactions that need energy to transport molecules and build the cellular structures necessary for all life processes
  • 3. Metabolic Processes
    • These reactions may be used to break down substances, synthesize others, and repair defective structures
    • Unneeded (or harmful) products of the reactions are eliminated as wastes
  • 4. The Chemistry of Life
    • Explore basis for biochemical processes
    • How atoms combine to form molecules used by animal and plant cells for daily bodily function (growth, maintenance, repair)
    • Understanding the chemistry of metabolism is key to understanding its many reactions and the products of those reactions
  • 5. 1.1 Atoms and Bonding
  • 6. Atoms and Elements
    • atom: smallest unit of matter involved in chemical reactions
      • nucleus = protons + neutrons
      • orbit/shell = electrons
  • 7. Atoms and Elements
  • 8. Atoms and Elements
    • atomic number = # of electrons = # of protons
    • atomic mass = # of protons + # of neutrons
            • OR
    • # of neutrons = atomic mass - atomic number
    * atomic mass is the average mass of all the elements naturally occurring isotopes
  • 9. Atoms and Elements
    • Atoms of the same element that contain different numbers of neutrons are called isotopes
    • Some isotopes are stable, whereas others are unstable and break down (decay)
    • The unstable isotopes are known as radioactive isotopes
  • 10. Electron Energy
    • Bohr model: electrons orbit the nucleus of an atom within energy levels, or shells
    • Electrons in the first shell (nearest the nucleus) have the lowest amount of potential energy
    • Electrons in the remaining shells have more potential energy
  • 11. Electron Energy
    • The chemical properties of atoms rely mostly on the number of electrons in the outermost occupied shell
    • Outermost shell = valence shell
    • Electrons in valence shell = valence electrons
    • Elements with 8 valence electrons are most stable (noble gases)
    • Other elements want to get to 8 electrons by donating or accepting electrons
  • 12. Ionic and Covalent Bonds
    • Most atoms can form chemical bonds with other atoms
    • These bonds are the forces that hold the atoms together in the form of compounds
    • Covalent Bonding: sharing of electrons between atoms
    • Ionic Bonding: transferring of one or more electrons from one atom to another
  • 13. Ionic Bonding sodium atom (Na) sodium ion (Na + ) 11 protons (+) 10 electrons (-) one (+) charge 11p 12n 11p 12n + electron given up
  • 14. Ionic Bonding 17p 18n 17p 18n electron accepted 17 protons (+) 18 electrons (-) one (-) charge chlorine atom (Cl) chlorine ion (Cl - ) -
  • 15. Ionic Bonding 11p 12n 17p 18n Na Cl +
  • 16. Ionic Bonding sodium chloride (NaCl) 11p 12n + 17p 18n -
  • 17. Ionic Bonding
    • If an atom loses an electron a positive ion is formed
    • A positively charged ion is called a cation
    • If an atom gains an electron a negative ion is formed
    • A negatively charged ion is called an anion
  • 18. Covalent Bonding Cl + Cl 17p 18n 17p 18n
  • 19. Covalent Bonding Cl-Cl chlorine gas (Cl 2 ) 17p 18n 17p 18n
  • 20. Covalent Bonding
    • Electron-dot or Lewis-dot diagrams show how atoms bond; only the valence electrons are drawn
    • Structural shorthand formulas can also be used; each line represents a pair of electrons shared by two atoms
    • Molecular formula only shows what type of atom is in the molecule
  • 21. Covalent Bonding
    • In a covalent bond, atoms share two valence electrons
    • Double covalent bonds involve the sharing of two pairs of shared valence electron
    • The shared electrons in covalent bonds belong exclusively to neither one nor the other atom
    How do we know if a bond is a covalent bond or an ionic bond?
  • 22. Electronegativity
    • A measure of the relative abilities of bonding atoms to attract electrons
    • The difference in electronegativity between two atoms helps predict the type of bond that will form (ionic, polar covalent, or covalent)
    • EN values can be obtained from the periodic table; elements near each other have similar EN values
    • Elements in the upper left have the largest EN values; elements in the lower left have the lowest EN values
  • 23. Electronegativity larger eletronegativity - smaller electronegativity = difference in electronegativity
  • 24. Electronegativity Examples: O 2 -> oxygen has an EN of 3.44 3.44 - 3.44 = 0 ∆ EN = 0, therefore each atom in an oxygen molecule has an equal attraction for the bonding pair of electrons. This is a covalent bond.
  • 25. Electronegativity Examples: HCl -> H has an EN of 2.20, Cl has an EN of 3.16 3.16 - 2.20 = 0.96 ∆ EN = 0.96, therefore chlorine has a slightly great pull on the bonding pair of electrons. This is a polar covalent bond.
  • 26. Electronegativity Examples: NaCl -> Na has an EN of 0.93, Cl has an EN of 3.16 3.16 - 0.93 = 2.23 ∆ EN = 2.23, therefore chlorine has a substantially greater pull on the bonding pair of electrons. This is an ionic bond.
  • 27. Electronegativity
    • A covalent bond is non-polar when both atoms share the electrons evenly, therefore there is no net positive or negative charge (∆EN is 0 to 0.7)
    • A covalent bond is polar when one atom has a slightly greater pull on the shared electrons creating a slightly negative atom and a slightly positive atom (∆EN is 0.7 to 1.7)
    • A bond is ionic when the ∆EN is greater than 1.7
  • 28. Polar Molecules
    • Due to the shape and difference in electronegativity of some molecules they can develop a slightly negative end (or pole) and a slightly positive pole
    • These are called polar molecules
    • Water is a common example of a polar molecule
    water (H 2 O)
  • 29. Hydrogen Bonding
    • Hydrogen bonds form between other molecules that contain hydrogen atoms bonded covalently to atoms of a much more electronegative element (eg. NH 3 , HF)
    • Hydrogen bonds form because of the polar nature of the molecule
    • A hydrogen bond is a force between molecules it is NOT a chemical bond within a molecule
  • 30. Hydrogen Bonding hydrogen bonding in water
  • 31. Solubility of Substances in Water
    • All cells depend on liquid water
    • Water comprises as much as 90% of a typical cell
    • Water is a perfect fluid environment through which other molecules can move and interact
    • Ionic compounds, like salts, dissolve readily in water
  • 32. Solubility of Substances in Water
    • Positively charged poles of water molecules attract anions in the salt
    • Negatively charged poles of water molecules attract cations in salt
    • These two attractions pull the sodium ions and chloride ions away from each other
    • The salt is now dissociated
  • 33. Solubility of Substances in Water
    • Compounds that interact with water — for example, by dissolving in it — are called hydrophilic
    • compounds that do not interact with water are called hydrophobic
    • Non-polar compounds are hydrophobic
    • Many protein molecules have hydrophobic regions in portions of their structure, and interactions of these regions with water cause the molecules to adopt specific shapes