Inorganic Chemistry/Biochemistry Notes
Inorganic Chemistry/Biochemistry Notes BIOL 1406 02
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This 6 page Class Notes was uploaded by locnaschek on Sunday September 4, 2016. The Class Notes belongs to BIOL 1406 02 at Lamar University taught by Dr. Randall Terry in Fall 2016. Since its upload, it has received 29 views. For similar materials see General Biology I (Majors) in Biology at Lamar University.
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Date Created: 09/04/16
Inorganic Chemistry Biochemistry is the chemistry of life. Cells are made up mostly of water, which is the basis of life. For example, the average human body would produce about eight pounds of ashes if cremated; the rest of the weight would evaporate as water vapor. Elements Critical to Life Macroelements present in living things in large quantities o The most important macroelements are CHNOPS, Ca, and K Oxygen (O), Carbon (C), Hydrogen (H), and Nitrogen (N) make up 96% of a human body’s weight Calcium (Ca), Phosphorus (P), Potassium (K), Sulfur (S), Sodium (Na), Chlorine (Cl), and Magnesium (Mg) make up 4% of a human body’s weight Microelements present in living things in small quantities but still essential o Boron (B), Chromium (Cr), Cobalt (Co), Copper (Cu), Fluorine (F), Iodine (I), Iron (Fe), Manganese (Mn), Molybdenum (Mo), Selenium (Se), Silicon (Si), Tin (Sn), Vanadium (V), and Zinc (Zn) Atomic Models There are three subatomic particles: o Protons (positive charge, located inside the nucleus) o Neutrons (neutral charge, located inside the nucleus) o Electrons (negative charge, located in the electron cloud surrounding the nucleus) An element in the Ground State has a net charge of zero, which means it has an equal number of protons and electrons The Mass number of an element refers to the number of protons plus the number of neutrons. Each of these subatomic particles has an atomic mass of 1 amu (atomic mass unit) and electrons have such a small atomic mass that they are not factored in. For example, the mass number of Helium (He) is 4 in its normal state (two protons + two neutrons = four amu). The Atomic number of an element is the number of protons in the nucleus. This never changes, and if it does you are looking at a different element. For example, Helium (He) has an atomic number of 2. If you see and atomic number of 3, you are now looking at Lithium (Li). o Mass number Atomic number = number of neutrons Electrons can and do change; these are responsible for the charge of the atom. They are located in energy levels, or shells, around the nucleus. The outermost level is called the valence shell. This outer energy level determines the reactivity of the atom. o Certain numbers of electrons fit into each shell (2 in the inner most, 8 in the second, 8 in the third, etc.) When the outermost shell of any particular atom is full, the atom is stable and not chemically reactive. For example, the far right column of the Periodic Table, the Noble Gases, all have full valence shells naturally, so all are stable. Atomic Bonding There are three kinds of atomic bonding. Covalent shares electrons The outer electron shell, the valence shell, wants to be full (usually eight) so sometimes it will share electrons with other atoms in order to achieve a full valence shell. Each column of a periodic table has the same amount of valence electrons and thus will react with a certain other column which can satisfy the desired full valence shell. Example 1: H—H is H , a 2ingle covalent bond symbolized by the singular dash. The one electron in each of their valence shells is shared so they each have two, creating a full outer shell. Example 2: O==O is O , a2double covalent bond symbolized by double dashes. They shared two of each of their six electrons, ending up with eight each and a full valence shell. o There are also polar covalent bonds; these form when atoms with different electronegativities react with one another causing one side of the molecule to be more positive and the other side to be more negative. Electronegativity is the tendency to attract electrons; it is unique to every element. Oxygen and Sulfur are amongst the most electronegative, which means that they have one of the strongest attractions of electrons. Example 1: H a2 O ar2 nonpolar. As both elements involved are the same, they have the same electronegativity and neither pulls in the other stronger. Example 2: in a water molecule, H O, the oxygen is more electronegative 2 than the hydrogen so the oxygen attracts the hydrogens more and causes the electrons to be shared unevenly. The hydrogen side becomes more positive and the oxygen side becomes more negative. If there are two different elements in a molecule, it will always be polar, even if it is only slightly. Ionic transferring electrons o As we have already discussed, the valence shell of atoms wants to be full, so sometimes electrons will be shared so onesided (as if one has 1% of the share and the other has 99% of the share) that they will essentially be transferred from one to another. o The transfer is caused by one element being so much more electronegative that it essentially takes the electron away from the other atom. o Example 1: Sodium (Na) has one valence electron in its ground state. Chlorine (Cl) has seven valence electrons. Cl is more electronegative, so it takes Na’s one and adds it to its own seven to get a full shell of eight. Na loses its one so if you move to the next shell in, it now has a full shell of eight as well. They bond to form sodium chloride (NaCl), also known as table salt. Na is a cation (positive charge) and Cl is an anion (negative charge). To find the charges: Na: neutral charge (0) minus one electron(1) equals (+1) charge [0(1)=0+1=+1] Cl: neutral charge (0) plus one electron(1) equals (1) charge [0+ (1)=01=1] Hydrogen a consequence of polarity, basically a covalent bond o Opposite charges attract and like charges repel o Water molecules are organized uniformly where the + and – sides are together o Hydrogen bonding is caused by polar covalent bonds; oxygen and hydrogen’s different electronegativities enable water molecules to bond to each other o Though hydrogen bonds are very weak, they are rarely found alone; when there is one, you are likely to find thousands, if not millions. When they are together they become tougher to break. It takes longer to boil a pot of water the more water you have in there because it takes so much energy to raise the temperature of the water to break the hydrogen bonds. A high water content inside cells helps the cell to maintain homeostasis and better resist temperature changes The strength of bonds can be determined by the amount of heat it takes to break the bond o Covalent bonds are usually stronger than ionic bonds, but they can overlap at times; both are much stronger than hydrogen bonds OxidationReduction (Redox) Reactions LEO the lion says GER: Loss of Electrons is Oxidation; Gain of Electrons is Reduction o What this means is that the atom that loses electron(s) is “oxidized” and the atom that gains electron(s) is “reduced” o Another term used is the “agent.” The oxidizing agent does the oxidizing (is reduced) The reducing agent does the reducing (is oxidized) o A way to describe this concept is a travel agent. The agent doesn’t travel; they help others travel. Likewise, the oxidizing agent doesn’t oxidize, it helps another atom to oxidize. o An oxidationreduction, or redox, reaction is a polar covalent reaction Example 1: In H O,2H is oxidized (loses an electron) and O is reduced (gains an electron). Example 2: In NaCl, Na is oxidized and Cl is reduced. Electron Transport Chains Electrons transferred in redox reactions start out with high energy and lose more and more energy as they are transferred. An Electron Transport Chain is a group of proteins (located in a membrane of an organelle in a eukaryotic cell such as a mitochondria or a chloroplast) organized in a particular way o The electronegativity gets higher from one protein to another (the electrons are pulled from one to another because of the increase in electronegativity) o At the same time, the energy level decreases from one protein to the next as energy is released, some of which is available to do work An Electron Transport Chain works by: o A molecule, or a “high energy electron donor,” gives two electrons to the first protein; these electrons are shuttled from protein one to protein two to protein three and on in an organized fashion o The diagram for an electron transport chain is shown in a descending staircase format to show the loss of energy. The energy released goes to cells. o The second molecule, the “terminal electron acceptor,” has a very high electronegativity which enables it to draw the electrons o This concept will be explained in more depth in a later chapter Role of Covalent and Hydrogen Bonding in Biochemistry A simple way to depict amino acids in a way that they can be identified is a beaded necklace type shape. Covalent bonds link together the separate amino acids in a protein, and they also hold the building blocks of the other macromolecules together. o Covalent bonds maintain the structural integrity of the macromolecule (for an image example of a protein, see slide 20) o Hydrogen bonds between the side chains of an amino acid determine the folding pattern of this necklace shape into a threedimensional blob The amino acids form hydrogen bonds with amino acids that are not next to them, these bonds are chosen carefully to ensure correct folding which is highly important in protein folding The bonds are weak which allows the blobshape to be more pliable o Shape and function are directly related; the function is determined by the shape
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