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Chapter 2 A quick review of Chemistry

Chapter 2 A quick review of Chemistry. What does chemistry have to do with biology?. All organisms are made of matter – In living things matter is constantly being rearranged through changes called chemical reactions.

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Chapter 2 A quick review of Chemistry

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  1. Chapter 2 A quick review of Chemistry What does chemistry have to do with biology?

  2. All organisms are made of matter – In living things matter is constantly being rearranged through changes called chemicalreactions Laws of chemistry -govern thestructure & function of alllivingthings. EXAMPLE of a chemical in the body: This crystal is uric acid, normally found in urine. It displays double refraction under a microscope with polarized light.

  3. Everything in the universe is made of matter • Matter- anything that occupies space and has mass. • Mass- the quantity of matter in an object -the total number of sub-atomic particles (electrons, protons and neutrons) in the object. **The chemical properties of matter depend on its organization at the following levels: Subatomic Atomic Molecular

  4. A. The Subatomic level(inside the atom) • Inside the Nucleus: • Proton- positively charged particle 2. Neutron-neutral, no electrical charge • Outside the nucleus 3. Electron –small, negatively charged particles The Total electrical charge of atom is zero

  5. * Atoms are composed mostly of empty space - between nucleus and the orbiting electrons • Electrons • high energy particles with very little mass. • move around outside the nucleus at very high speeds in areas called orbitals. • An orbital- is a region around the nucleus that indicates the probable location of an electron.

  6. Isotopes • All atoms of an element have the same # of protons. • All atoms of an element do not necessarily have the same # of neutrons. • Isotopes are atoms of the same element with a different # of neutrons. • The additional neutrons change the mass of the atom. • Most elements are a mixture of isotopes.

  7. Atomic level-atoms of each element • Element -A substance that cannot be broken down by chemical means. Elements are defined by their number of protons (Atomic Number) Example- an atom of Hydrogen, (H -1)Oxygen (O -8), Carbon (C-6) *The number of electrons in the outer most energy level determine bonding properties

  8. Elements in Living things- • Approx. 25 elements occur in living things. • 6 elements of critical importance: • Carbon – Hydrogen – Nitrogen – Oxygen – Phosphorous - Sulfur • “CHNOPS” • Found in the 4 major groups of ORGANIC molecules: • Carbohydrates - Lipids • Proteins -Nucleic Acids

  9. In the Human BodyNote the amounts of “CHNOPS” plus Ca

  10. The Periodic table

  11. Groups or “Families” • vertical columns in the periodic table • -share characteristics & traits • -groups have similar configurations of the outermost electron shells of their atoms. • Example: Group 14 is the carbon group- Each of the elements in this group has 4 electrons in its outer energy level. • That leaves 4 empty spotsfor additional electrons- so they form 4 chemical bonds

  12. C. Molecular level • Compound-a pure substance made up of atoms of two or more elements **Most elements undergo chemical reactions- combine in ways that arrange their atoms so the outer electron shell is full & stable **Compounds have specific formulas Molecule-the simplest part of a substance that retains all of the properties of the substance and can exist in a free state

  13. 2.Chemical bonds –atoms want to form stable molecules - Covalent Bonds – some elements are more stable when sharing electrons - Ionic Bonds- some elements really want to lose or gain an electron- forming cation (+) or anions (-)

  14. Covalent Bonds Formed when 2 or more atoms share one or more pairs of electrons *Example- water

  15. Forming ions

  16. Ionic Bonds Form when a positively and a negatively charged particle called IONS attract each other These positive and negative IONS form when an atom loses or gains an electron to become “stable” and have a complete # of electrons in the outer shell Example- LiF

  17. Energy transfer in Living Things • Definition of Energy- the ability to do work or cause change • Free Energy –the energy in a system that is available for work For example, in the cell, it is the amount of energy available to fuel cell processes. Energy can be converted- breakfast food (chemical E) -to thermal E & mechanical E

  18. C. Chemical Reactions-many complex chemical reactions occur in living things every second • Reactant: substances that enter a chemical reaction • Product: substances produced by chemical reactions

  19. Example: Chemical reaction in body Reactants & Products CO2 + H20 == H2CO3 Reactants Products Reaction is reversible- shows the reaction of water and carbon dioxide to carbonic acid in blood

  20. Which are the reactants & which are products in each reaction. Is this reaction reversible? Explain. Photosynthesis: • 6CO2 + 6H2O + Energy (sun) C6H12O6 + 6O2 • Cellular metabolism: • C6H12O6+ 6O26CO2 + 6H2O + Energy(ATP)

  21. D. Energyfor reactions 1.Activation Energy-The amount of energy needed to start a chemical reaction • Catalysts &Enzymes– chemical substances which are able to lower the amount of activation energy needed for a reaction. • Enzymes-are biological catalysts- they lower the activation Energy needed for chemical reactions in living things. Why are enzymes so important?

  22. Features of a catalyzed reaction Lower activation E needed for reaction w/ catalyst

  23. Enzymes and chemical reactions in the body • Lactase is one of many hard working enzymes. It breaks up the chemical bond holding the sugar lactose together. • What happens if you don’t have enough of this enzyme in your body?

  24. III. Water & Solutions • A. Polarity • Water is a polar molecule • Polar molecules -have an uneven distribution of charge, though they have a net zero charge. • Electrons in a water molecule are shared unevenly between hydrogen and oxygen

  25. Uneven electron sharing in the covalent bond in water http://www.ciese.org/curriculum/waterproj/images/watermolecule.jpg

  26. Solubility of Water • The polarity of water makes it effective at dissolving other polar substances such as sugars, ionic compounds, and some proteins. • Water is considered the Universal Solvent because so many things can dissolve in it. • (non-water soluble means it is not polar- like oil, fats & other lipids)

  27. Hydrogen Bond • is the force of attraction between a hydrogen molecule with a partial positive charge and another atom or molecule with a partial or full negative charge

  28. III. C. Solutions Definitions 1. Solution-a mixture in which one or more substances are uniformly distributed in another substance. • Solute– the substance dissolved in the solvent. • 3. Solvent– the substance in which the solute is dissolved. Example: salt is the solute, and water is the solvent

  29. 4. Concentration- the measurement of the amount of solute in a fixed amount of solution. Example- a 5 % sugar solution contains 5 grams of sugar + enough water to make 100 ml of solution.**Concentration increases with amount of Dissolved solute

  30. 5. Saturated Solution- solution in which no more solute can dissolve. 6. Aqueous Solution- water is the solvent What aqueous solutions can you think of that are important to living things?

  31. B. Acids & Bases Water forms ions: H2O H+ + OH- The H+ then combines with another water to make H3O+ Called a hydronium ion The OH- is called a Hydroxide ion 1. Acid – Has more H+ (Hydronium ions) 2. Base – Has more OH- (Hydroxide ions)

  32. pH scale • The pH scale measure the amount of H+ (acid) or OH- (base) in a solution

  33. A change of 1 pH unit shows a tenfold increase in H+ or H3O+ ions. • For Example -a pH of 5 has 10 times more H+ than pH 6 What Ph is an acid? A base? What pH is Neutral?

  34. Buffers- chemicals that neutralize small amounts of either acids or bases. • If you put both an acid and a base into the same container of water, they tend to cancel out the effects of one another. • For example, if both HCl and NaOH are placed in water, the Na+ and Cl- ions combine to form NaCl (table salt), and the H+ and OH- ions combine to form H2O (water). • Buffers in the body are very important- you have very acidic stomach fluid and basic blood.

  35. Buffering in the body This figure shows the major organs that help control the blood concentrations of CO2 and HCO3-, and thus help control the pH of the blood. Removing CO2 from the blood helps increase the pH. Removing HCO3- from the blood helps lower the pH.

  36. Indicators are usually organic compounds which change color in relation to the concentration of hydrogen ions. pH indicators • pH test strips ( Litmus papers) are quick and work by color change

  37. Tasting acids and bases(Note: NEVER taste an unknown solution)Acids -Vinegar and citrus solutions have a sour taste. Bases - Baking soda & unsweetened chocolate have a bitter taste. Sour Taste Receptors Bitter Taste Receptors The sides of the tongue have adapted to taste sour substances. The taste buds at the back of the tongue have adapted to taste bitter substances. Other receptors are specialized for tasting sweet and salty.

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