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The Chemistry of Life

The Chemistry of Life. Elements. Only 90 naturally occurring elements and only 25 of those are essential to living organisms. Four elements make up more than 96% of the human body: Carbon, hydrogen, oxygen, and nitrogen. (trace amounts of P and S=CHONPS). Periodic Table.

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The Chemistry of Life

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  1. The Chemistry of Life

  2. Elements • Only 90 naturally occurring elements and only 25 of those are essential to living organisms. • Four elements make up more than 96% of the human body: Carbon, hydrogen, oxygen, and nitrogen. (trace amounts of P and S=CHONPS)

  3. Periodic Table • Rows are called periods and are arranged by the number of orbits they contain. (2, 8, 8, 16) • Columns are called families and like families they have similar bonding qualities. • You can take any element and tell how many orbits it will have and how many valence electrons it will have based on it’s row and column. • A zig-zag line separates the metals from the non-metals.

  4. Atoms • An element has only one type of atom. • From the Greek word “atomos”, which means unable to be cut • A million atoms placed side by side would make a row only 1 centimeter long • Contain subatomic particles: • Protons, neutrons & electrons

  5. Nucleus • At the center of the atom • Contain positively charged protons and neutral neutrons that are held together by strong forces. • Protons & neutrons are about the same size and make up all the mass of the atom.

  6. Electrons • Negatively charged electrons orbit around the nucleus much like the earth orbits around the sun • They are much smaller than protons and neutrons: about 1/1840 the mass of a proton, so they do not add to the atomic mass. • Electrons are constantly in motion. • They are attracted to the positively charged nucleus.

  7. The overall charge of an atom is NEUTRAL • Atoms have the same number of positively charged protons as they do negatively charged electrons. • Even though their masses are different, their charges still count equally. • 2 protons(+) and 2 electrons(-) would equal a neutral charge.

  8. Isotopes • An element is identified by its number of protons, called the atomic number. • The atomic mass is the number of protons plus the number of neutrons in the nucleus. • An element can have different numbers of neutrons, but not different number of protons or it would be a different element. • These different versions of the same element are called isotopes. • Isotopes will have the same chemical properties.

  9. Atomic mass • The atomic mass listed on the periodic table is the average of the isotopes of the element found on earth multiplied by the percentage at which they are found. • The mass is calculated this way in order to more accurately represent the way the element is found in nature.

  10. Radioactive Isotopes Some isotopes of certain elements are unstable and will break down over a period of time. This means they will lose neutrons at a constant rate and change into other isotopes of the same element. Since they will not lose protons, they are still the same element, just different versions of that element.

  11. Uses of radioactive isotopes • Carbon dating using Carbon 14, which breaks down to Carbon 13 & Carbon 12. • As isotopes break down, they give off radiation and can be used to treat cancer, as tracers to follow the movement of substances in organisms, and to kill bacteria that cause food to spoil.

  12. Compounds • When 2 or more elements form a stable union, it is called a compound or a molecule. • A compound is a combination in definite proportions such as H2O is always water, H2O2 is hydrogen peroxide. • A molecule is just the smallest form of a compound and the two terms are used interchangeably. • They can be same element or different elements.

  13. Every molecule has unique characteristics that differentiate it from every other molecule. • Even a small change can make a big difference • Glucose and Fructose are both C6H12O6, but just where their atoms are located make them different and we can even taste the difference.

  14. Molecules are important in life • DNA and RNA are distinguished only by the removal of one oxygen atom. • DNA is double stranded and more stable than the single stranded RNA

  15. Chemical Reactions • Compounds and molecules are created when bonds are formed between elements. • REACTANTS ------- PRODUCTS • Na + I2  2 Na I • Balancing Equations

  16. Bonds • Chemical bonds have different strengths. • Ionic Bonds are strong and form crystals, which take a lot of energy to break apart, but they easily dissolve in water. • Covalent bonds are the most common. • Hydrogen bonds are the weakest, but are strong when there are many of them together. Bond Animation

  17. Ionic Bonds • Ionic bonds occur between metals and non-metals on the periodic table. • The valence electrons are transferred from the metal to the nonmetal. • An ion is a molecule that creates ions and the positive and negative charges created draw the atoms very close together. • The atoms involved have a strong need to fill or empty an orbit.

  18. Table Salt Crystal

  19. Ionic Bonding • Turn to your periodic table and examine the two columns headed by Li (ignore hydrogen), Be. • These columns provide most (not all) of the positive partners involved in ionic bonding that a high school kid will be held responsible for.

  20. Ionic Bonds • The first column (called the alkali metals) has Li, Na, K, Rb, Cs, and Fr • All these guys go +1 in ionic bonding. • The second column (called the alkaline earth metals) has Be, Mg, Ca, Sr, Ba and Ra. All go +2 in ionic bonding

  21. Ionic Bonds • Look to the column headed by F and below it, you'll see Cl, Br, I and At • These elements will all gain one electron in ionic bonding and will therefore be negative one. • The next column to the left is headed by O. The most common examples used from this column are O and S. Se and Te get used sparingly • Gaining two electrons makes these atoms become a negative two charge in ionic bonding.

  22. Electronegativity • Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. • The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7. • If two elements have an electronegativity greater than 1.9, they will form an ionic bond.

  23. Electronegativity • Metals have low electronegativities (less than 2.0), while non-metals have high electronegativities (above 2.0) • Thus, nonmetals from the right side of the chart will ionically bond with metals from the left side of the chart. • The only way you would know electronegativity is if you were given a chart.

  24. Covalent Bonds • A covalent bond forms when electrons are shared instead of transferred between atoms. • Atoms can share one electron for a single covalent bond, and even 6 electrons for a triple bond. • The water molecule is a great example of covalent bonding.

  25. H2O

  26. Methane CH4

  27. Hydrogen Bonds • The positive hydrogen in polar molecules forms weak bonds with the negative poles of other polar molecules. • These weak bonds are called HYDROGEN BONDS and are very important in biological molecules. • Hydrogen bonds allow DNA to be stable, yet easily pulled apart to be read for instructions.

  28. van der Waals forces • Intermolecular forces that hold molecules together based on their slight polar charges. • Little fibers on the foot of a Gecko allows it to climb walls using the van der Waals forces.

  29. Properties of Water

  30. Biologically Important • ¾ of the earth’s surface is covered with water. • Water makes up about 70% of the mass of the human body. • Life as we know it could not exist without water.

  31. Types of Chemical Bonds

  32. Shape of the water molecule:

  33. Electrons are not shared equally. • Polar Covalent with partial positive and partial negative charges.

  34. Hydrogen bonds form between water molecules. • Individual bonds are relatively weak • But multiple bonding holds molecules close together making it overall fairly strong.

  35. Hydrogen Bonds • An intermolecular force of attraction • A relatively weak bond (about 1/10 the strength of a covalent bond). • Occurs between a hydrogen atom and an oxygen, nitrogen, fluorine.

  36. Thermal Properties of Water • Unusually high melting and boiling points compared to other molecules of similar mass. • Resists phase change • Indication of how strongly the molecules of water are attracted to each other.

  37. High heat capacity • calorie:The amount of energy required to raise a defined amount of a substance by one degree.

  38. Effect on Body Temperature • Large change in surrounding temperature produces a small change in temperature of water. • Helps maintain a fairly constant body temperature (37o C) • Essential for normal biochemical processes

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