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  1. THE CHEMISTRY OF LIFE UNIT 1 - Chapters 2,3

  2. S Ag Hg 87 Tc - Technitium Francium 92 Natural Elements • Silver, Sulfur and Mercury are examples of naturally occurring elements • Others are laboratory synthesized radioactive elements - some of these decay rapidly. Technitium, a radioactive silver-gray metal was first to be synthesized. Plutonium, Promethium, Francium are other examples. Some of these man made elements (like Francium) actually do occur in nature in extremely minute amounts

  3. Goiter – Thyroid enlargement 25 Elements Needed for Life • 25 of the 92 elements are found in all life forms • The 4 most common elements make up 96% of a cell. They are: Hydrogen (H), Oxygen (O), Carbon (C), Nitrogen (N) • Others such as Phosphorus (P), Sulphur (S), Potassium (K), Calcium (Ca), etc. account for the remaining 4% • Trace elements are those required in minute quantities. Eg. Boron, Iodine, Iron, chromium, zinc, manganese, selenium, silicon, tin, vanadium, molybdenum, cobalt, copper and flourine.

  4. Atoms and subatomic particles • Atoms are the smallest units of matter that have the properties of the element they represent • An atom can be split into many different subatomic particles, but only three are stable enough to have been studied for many decades: Protons, Neutron and Electrons • Protons – positively charged, have a mass of about 1 dalton • Neutrons – electrically neutral, mass close to 1 dalton • Electrons – negatively charged; their mass is about 1/1800 of protons and neutrons, so it can be ignored • So the mass of an atom = number of protons + number of neutrons • The atomic number = the number of protons

  5. Who the heck is Avogadro and why do I need his number? • Atoms, molecules too small – so how do we measure them for experiments? • Scientists tried to determine how many particles, in a 1 cubic centimeter area. • Loschmidt, Perrin and Einstein had more of a role to play in Avogadro’s number than Avogadro himself, although he came up with early, vague numbers.

  6. Avogadro's Number • The atomic mass of hydrogen is 1 g/mol. There are 6 x 1023 atoms in a mole. So the mass of one atom of hydrogen = 1g / 6 x 1023 =1.7 x 10-24 g (which is close to 1 dalton) In this case, the H atom only has a proton, no neutron and the electron is so light, its mass does not matter. Therefore, the mass of each proton and each neutron is about 1 dalton or 1.7 x 10-24 g

  7. Atomic mass, molecular weight, moles and all that stuff • What is the atomic mass per mole of Helium? (What is the weight of a dozen chicken eggs?) 4 grams per mole (60 grams per dozen) • How many atoms in a mole of Helium? (How many eggs in a dozen chicken eggs?) 6 x 1023 atoms(12 eggs) • How much does each atom of Helium weigh? (How much does each chicken egg weigh?) 2 protons + 2 neutrons = 4 g / 6 x 1023 = 6.7 x 10-24 g per atom of Helium (60 g / 12 = 5 g per chicken egg)

  8. Fundamental Particles The nucleus of an atom normally contains two types of particle: the proton and the neutron. From what you have learned before, it may seem that protons and neutrons are fundamental particles. This would be a nice simple picture of matter: three particles that make up the Universe - protons, neutrons and electrons. However, protons and neutrons are only two particles in a large family called hadrons. Unlike the electron, hadrons are not fundamental - they are made up of even smaller particles called quarks. Quarks are fundamental. They make up one family of fundamental particles. The other family is the leptons (the electron's family). Although the quarks are a family of fundamental particles, they never exist on their own. They are only ever found as combinations in protons, neutrons and the other hadrons. Protons and neutrons are made of two types of quark. These quarks are said to have different flavors: up and down. These up and down quarks are the only quarks that are found in normal matter and they are known as first generation quarks. A proton is made from two up quarks and a down quark. A neutron is made from two down quarks and an up quark. Table 1 shows the properties of these quarks and how they combine to give the charges of protons and neutrons. -1/3 +2/3 +2/3 =+1 -1/3 -1/3 +2/3 =0

  9. Isotopes • Atoms of an element can exist in alternate forms – some can be “heavier” and some “lighter” • A sample of any element contains a mixture of these versions, but the original atom (represented in the periodic table) is the most abundant • These alternate versions are called isotopes • Basically, isotopes have varying number of neutrons but the same number of protons (So they have different atomic masses, but the same atomic number) • Isotopes containing fewer neutrons than the most abundant atom of the element are called heavy isotopes and those that have fewer neutrons are called light isotopes

  10. Beta particle or electron mass is Too small to affect atomic mass. It Stays in orbit around nitrogen nucleus. Radioactive Isotopes • Usually unstable • Nuclei of radioactive atoms have a tendency to “decay” • Decay means to give off particles and energy • For example, in beta decay, a nucleus throws out a fast moving electron (a.k.a a beta particle) as a neutron’s down quark changes into an up quark making it a proton. • An example of beta emission is the decay of carbon-14 into nitrogen-14. The equation for the decay is:

  11. What does radiation do to us? • If radiation is passing through your cells, which are primarily made of water, it will causes electrons to leave some of the atoms in the cell. These atoms then will have a charge (an ion) and can go on to react with other atoms in the cell, causing damage. An example of this would be if a gamma ray passes through a cell, the water molecules near the DNA might be ionized, then the ions might react with the DNA causing a break in it. • With enough damage, the cell might die. Your body though has trillions of cells and repairs most of the damage quickly. Radiation amounts which this cell death might cause a physical problem like radiation sickness can only happen during major accidents and so are extremely rare.

  12. Electrons • Spin in “energy levels” or shells, at ground state (unexcited state) • The farther the shell is from the nucleus, the more potential energy of the electrons (because of their attraction to the protons in the nucleus) • An electron can move to a higher shell by absorbing energy (excitation by light); or a lower shell losing energy (releasing heat and light in the form of photons) • They may or may not return to ground state

  13. Filling Atomic Shells • In general, atoms are most stable when they have 8 electrons in their outer-most shell. (Octet means 8.) The exception is the first shell which is most stable with TWO electrons. • Electrons having opposite spins are said to be "paired" electrons • The electrons in the outer shell are called the valence electrons 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 S subshells carry  a maximum of 2 electrons and have 1 orbital. P subshells carry a maximum of 6 electrons and have 3 orbital. D subshells carry a maximum of 10 electrons and have 5 orbital. F subshells carry a maximum of 14 electrons and have 7 orbital. Therefore each orbital carries 2 electrons. An orbitalis a region of space where there is a 95% chance of finding an electron.

  14. Filling Electron Shells, cont’d.

  15. Why fill 4S first then 3d?

  16. Chemical Bonds When atoms complete their valence shell by either sharing or transferring unpaired valence electrons, they tend to stay close together – this is a chemical bond

  17. Covalent Bonds Sharing of valence electrons by atoms – Extremely strong bonds a. Non-polar – electrons shared equally – hydrophobic - single, double and triple bonds b. Polar – electrons spend more time closer to the more electronegative atom - hydrophilic Electronegativity is the tendency of an atom to attract electrons. The more one atom attracts electrons, the more electronegative it is e.g. O and N are very electronegative and their bonds with hydrogen will be polarized.

  18. Ammonium ion Co-ordinate (dative covalent) bonding • In the formation of a simple covalent bond, each atom supplies one electron to the bond - but that doesn't always have to be the case. A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.

  19. Electronegativity As the distance of the outermost electrons from the nucleus decreases, electronegativity increases. An easy way to remember is: electronegativities increase from the bottom left of the periodic table (francium, Fr) to the top right (fluorine, F), not including the noble gases. Compounds containing atoms with a large electronegativity difference form ionic bonds. This type of bonding can be described by assuming that the atom from the left side of the periodic table (low electronegativity, outermost electrons loosely held) gives up all of its valence electrons to the right-side atom (high electronegativity, outermost electrons tightly held). Therefore, a pair of oppositely charged ions is created, with an attractive electrostatic force between them that is the basis of an ionic bond.

  20. Electronegativity Valuesfor Some Elements

  21. Ionic Bonds • Compounds containing atoms with a large electronegativity difference form ionic bonds. • Atoms from the left side of the periodic table (low electronegativity, outermost electrons loosely held) gives up all of its valence electrons to the right-side atom (high electronegativity, outermost electrons tightly held). Therefore, a pair of oppositely charged ions is created, with an attractive electrostatic force between them that is the basis of an ionic bond.

  22. Difference between Covalent and Ionic bonds

  23. Weak bonds and attractions Hydrogen bond – when a hydrogen atom in a polar molecule is attracted to another electronegative atom Van der Waals forces – weak, transient attractions that form between atoms and molecules due to asymmetrical distribution of electrons


  25. Water In liquid water at 37 ˚ C, each water molecule has hydrogen bonds with 4 other water molecules. These weak bonds constantly break and form with other water molecules nearby – this gives water its fluidity.

  26. Properties of Water • Cohesion – the “sticking together” of water molecules to each other due to hydrogen bonds. Cohesion of water molecules creates High Surface Tension – a measure of how difficult it is to “break” the surface of water • Adhesion – the clinging of water molecules to molecules of another substance • High Specific Heat – Water is able to absorb or lose more heat than other substance, before its temperature changes. For example, iron has a lower specific heat index than water (it heats and cools faster than water) The specific heat of water is 1 calorie per gram per degree Celsius. • Evaporative Cooling – water molecules on the surface of water, are warmest and fastest. They tend to escape as gas (water vapor), taking away the heat. The next layer of molecules is cooler now, but will eventually heat up like the previous layer. (Dry day vs. muggy day)

  27. Water Transport in plants • Adhesion, cohesion and transpiration allow the supply of water to the tallest of trees, by forming a continuous column of water

  28. Properties of Water, cont’d. • In water vapor, molecules are too far apart to form hydrogen bonds • In liquid water, hydrogen bonds form and break constantly • In ice, all water molecules form long-lasting hydrogen bonds with 4 other molecules, creating a lattice • In ice, water molecules are farther apart from each other than in liquid, so there are fewer molecules in ice than in an equal volume of liquid water • Ice is therefore less dense than liquid water • Water at 4˚ C, is at its densest, since the molecules are close to each other and have formed longer lasting hydrogen bonds – this is because they are less agitated

  29. Evaporation Water molecules in liquid phase

  30. Water – the solvent of life • The polarity of a water molecules helps it dissolve ionic and hydrophilic substance easily

  31. Hydrophobic vs. Hydrophilic

  32. Hydroxide ion Hydronium or hydrogen ion Water molecules can dissociate • Sometimes a hydrogen in one water molecules can move to another water molecule • In the process however, it leaves behind its single electron with the old water molecules and only donates a proton to its new water molecule • The previous water molecules has a (-) charge and is called a hydroxide ion • The new water molecule has a (+) charge and is called a hydronium ion or just a hydrogen ion • This dissociation is reversible

  33. Ion Concentrations • In pure water, only 1 in 554,000,000 molecules of water is dissociated • In pure water, the concentration of hydrogen and hydroxide ions is equal and is 10-7 M each (This means that in 1 liter of pure water, there is only 0.00000001 mole of either H+ or OH-) • H+ or OH- concentrations can be changed by adding certain substances

  34. Acids and Bases • Acids are substance that “donate” H+ (protons) to an environment like water • Bases are substances that either donate OH- ions to an environment, or pick-up H+ (protons) from an environment like water

  35. Other properties • Acids taste sour, are corrosive to metals, change litmus (a dye extracted from lichens) red, and become less acidic when mixed with bases • Bases feel slippery, change litmus blue, and become less basic when mixed with acids

  36. Chemical reaction uses a single arrow implying the reaction is one way Strong Acids

  37. Chemical equation uses double arrows that imply the reaction is reversible Weak Acids

  38. Weak and Strong Bases • Similar to weak and strong acids, strong bases dissociated completely and the reaction has a single arrow whereas weak bases do not dissociate and the reaction is reversible (double arrows) NaOH  Na + + OH – Acid + Base = Salt + Water 2HCl + 2NaOH  2NaCl + 2H2O NH3 + H+ NH4+ • Weak acids and bases make excellent buffers. Buffers are chemicals that keep the acidity or alkalinity of an environment fairly stable, by picking up excess protons or adding needed protons (H+). Carbonic acid (H2CO3) is an excellent buffer in living cells

  39. The pH Scale • pH stands for Potential of Hydrogen • It is a measure of the concentration of H+ in a solution • The product of hydrogen and hydroxide ions in any given solution always equals 10-14 M • The pH of a solution is the negative log of the Hydrogen ion concentration or – log [H+] • In pure water, the [H+] is 10-7 M, so – log [10-7] = -(-7) = 7 The pH of pure water is 7, or neutral

  40. Increasing [H+] Decreasing [OH-] [H+] and [OH-] equal or 10-7 M each Increasing [OH-] Decreasing [H+]

  41. Acid Rain and the Environment

  42. So what does acid rain do?

  43. End of Chapters 1 and 2