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AP Biology Chapter 2 The chemical context of Life 1 Chemical elements and Compounds The atomic structure determines the behavior of elements. Atoms combine chemically to form compounds and molecules. A molecules biological function is related to its shape. Chemical reactions break chemical bonds. Organisms are composed of matter. Life requires about 25 chemical elements 96% Living Matter Carbon Hydrogen Oxygen Nitrogen 4 % Remaining Phosphorus Sulfur Calcium Potassium Trace elements are in organisms in minute quantities. Iron and iodine are examples of these. Atoms and Molecules The nucleus of the atoms contains protons and neutrons. ELECTRON Mass = 9.11 x1031 kgms Charge = -1 PROTON Mass = 1.67 x1027 kgms Charge = +1 NEUTRON Mass = 1.67 x1027 kgms Charge = 0 The nucleus of an atom has a high density compared to the overall atom. The nucleus is very small compared to the overall atom. The arrangement of the electrons determines the atoms chemical properties. Atoms can have isotopes. Isotopes have the same number of protons and electrons but a different number of neutrons. 2 Atomic Number is the number of electrons and protons. The total mass of the atom is Atomic Weight. 23 11 The element name is sodium Na 23 is the mass number 11 is the atomic number Molecules and Bonding Forces that hold atoms together are chemical bonds. Covalent bonds share electrons ---These include compounds such as water, alcohol, and most organics. Most compounds that are covalently bonded form within the “p” block elements. Organic compounds are covalent in nature and usually consist of carbon, hydrogen, oxygen, and nitrogen. Ionic Bonds exchange electrons ----These include inorganic salts like NaCl or NaOH 3 Binary compounds are formed by the action of metals reacting with non-metals and mutually transferring electrons. This results in the formation of cations and anions that mutually attract each other. This attraction causes a change in energy resulting in the formation of a binary or ionic compound. Energy is also a by-product. Even though a solid compound is formed, for the electrons to transfer a phase change takes place resulting in an energy change. This energy is Latice Energy or the change in energy that takes place when separated gaseous ions form together to make an ionic solid. The following reaction can happen: M+(g) + X-(g) MX(s) + ENERGY Reactions can happen in other compounds. Chemical Formula H 2O Str uct ura l For mu la Consider the ionic reaction: Na Na e Sodium looses an electron (OXIDATION) Element Consider the ionic reaction: Cl e Cl (REDUCTION) Element Cation or charged particle Chlorine gains an electron Anion or charged particle 4 Consider Avogadro’s Hypothesis for chemical reactions Polyatomic ions consists of several elements or atoms: NH 4Cl Ca3 PO4 AmmoniumChloride Calcium Phosphate morethantwo elements Group I Elements 1) A +1 charge 2) Also calles the Alkali Metals. 3) React with water to form an hydroxide. 4) These elements do not occur in nature but in the form of ores. 5) Ionize very easily. Group II Elements 1) +2 Charge 2) Also called the Alkaline Earth Metals 3) React slowly with water to produce a metal oxide. 4) These elements also do not occur in nature but rather in ores. Transition Elements 1) Mostly alloy with other elements especially with steel to make it tougher, ware resistant, and more durable. 2) Several of these elements have multi oxidation states. 3) The salts of these elements are very colorful. 4) Most of these elements conduct electricity. Halogen Gases 1) These elements are all diatomic. 5 2) These elements ionize very well especially with Group I elements and somewhat with Group II elements. 3) All have a –1 charge. Noble Gases 1) All are non-reactive. 2) All have a “0” charge. 3) Occur in nature as a mixture with other gases. Lanthanides and Actinides These elements are for the most part radioactive and used for military purposes or medical applications. Atomic Theory Atomic orbitals: Electrons can only occupy so-called atomic orbitals with well defined energy levels corresponding to the principal quantum number, n. The lowest level will have n = 1, the next n = 2, and so on. The maximum number of electrons which can "fit" into these energy levels is given by the formula 2n2, where n is the principal quantum number. So, the first energy level will accommodate a maximum of 2 electrons (2 x 1 x 1), the second level 8 (2 x 2 x 2), the third 18 (2 x 3 x 3) and so on. Within a main energy level with a specific value of n, there may be subsidiary levels which are designated s, p, d and f in increasing order of their energy values. The diagram on the left shows the relative energies of the atomic orbitals for principal quantum numbers n = 1 to n = 4. Note that the level with n = 1 is limited to a single s orbital. The level n = 2 has both s and p orbitals, while at the level n = 3, we have d orbitals. At the level n = 4, we have, in addition to s and p orbitals, d and f orbitals, which are not shown in the diagram. 6 The Aufbau Process and Energy Levels It is quite easy to establish this distribution for a given atom, by making use of simple rules. The best way to learn these rules is to see how they are applied in practice, in the so-called AUFBAU (building-up) process, for atoms in the ground state (that is, with all of its electrons in the lowest possible energy levels). Electrons must always enter the first available orbital of lowest energy. The first element, hydrogen, only has one electron, and so this electron must enter the 1s orbital. The electronic configuration of hydrogen in the ground state must therefore be: H 1s1 Pauli's exclusion principle must now be applied - the next electron to enter the 1s orbital must have a spin opposite to the spin of the electron which is already there. This completes the occupancy of the 1s orbital. The electronic structure of helium is: He 1s2. Consider the example The next two electrons enter the 2s orbitals in the same way. This leads us first to the element lithium, with 3 electrons (1s2 2s1) and then, beryllium, with 4 electrons, and an electronic structure 1s2 2s2. 7 The next element, boron, has 5 electrons. The first four of these fill the 1s and 2s levels, while the fifth enters the next level of lowest energy, which is a 2p orbital. The electronic structure of boron is: 1s22s2 2p1. We now apply Hund's rule: fill a set of orbitals of equivalent energy (the 2p orbitals in this case) in such a way that as many electrons as possible remain unpaired. The 6th electron enters a vacant 2p orbital rater than pairing with an unpaired electron. This element is carbon: 1s22s2 2p2. In this way, the next electron enters the vacant 2p orbital, giving nitrogen (1s22s2 2p3), oxygen (1s22s2 2p4), fluorine (1s22s2 2p5) and neon (1s22s2 2p6). s-Orbitals An orbital may be considered as a region in space to which electrons have access. Heisenberg's uncertainty principle states that one cannot simultaneously determine with certainty both the position and the momentum of an electron. So, an orbital may be represented as a region where there is a given probability of finding an electron at any given time. 8 The diagram on the right shows how the probability density, (P), of a 1s electron varies with its distance (r) from the nucleus. This so-called electron density of an electron in a 1s state is the same in all directions. A 1s orbital may be considered as a spherical distribution of negative charge, which becomes more diffuse as the distance from the nucleus increases. One tries to depict this graphically, as shown on the right. This is often called an electron cloud. 2s orbitals are similar in shape, but the radius of the electron cloud is larger. p-Orbitals p-Orbitals consist of electron clouds which look like the diagram on the right, with a node at the nucleus of the atom. There are three such orbitals, oriented at right angles to one another, as shown in the diagram below: 9 Combining all three p orbitals on the same set of axes, as shown on the left, gives an idea (but only a rough idea!) of the space available to an electron in the p-state. Remember that each orbital can only hold a maximum of 2 electrons (which must have opposite spins). The complete set of the three orbitals px, py and pz can therefore altogether accommodate a maximum of 6 electrons. Ions: Atoms are electrically neutral, but they may either lose or gain electrons, becoming charged ions in the process. If an atom loses one or more electrons, it will become positively charged and the resulting ion is called a cation (see a table of common cations): On the other hand, by gaining one or more electrons, an atom becomes negatively charged, and is then called an anion (see a table of common anions): Let's see what happens to the electron configuration of a sodium atom when it loses an electron to become the cation 10 The change in electron configuration is described by But 1s2 2s2 2p6 is the electron configuration of the inert gas neon (Ne), which, like others in that group, shows remarkable stability. What happens with chlorine? It gains an electron to form the chloride anion The change in electron configuration is described by But 1s2 2s2 2p6 3s2 3p6 is the electron configuration of the element 18, which is the inert gas argon (Ar). In general, metallic elements tend to lose electrons to form a stable octet of electrons. The more reactive non-metals such as oxygen and chlorine achieve this octet by gaining electrons. Light energy as from the sun can excite electrons to form new compounds. This concept is demonstrated in photosynthesis. Some common Cations Charge = +1 Charge = +2 Charge = +3 Hydrogen H+ Magnesium Mg2+ Aluminium Al3+ Lithium Li+ Calcium Ca2+ Iron(III) Fe3+ Sodium Na+ Barium Ba2+ Potassium K+ Zinc Zn2+ Silver Ag+ Iron(II) Fe2+ Copper(I) Cu+ Copper(II) Cu2+ Ammonium NH4+ Lead(II) Pb2+ 11 Some common anions: Charge = -1 Charge = -2 Charge = -3 Fluoride FChloride ClBromide BrIodide IHydroxide OH- Oxide O2- Hydrogen sulphide(1) HS- Sulphide S2- Nitrite NO2Nitrate NO3Hydrogen carbonate(2) HCO32- Carbonate CO22- Hydrogen sulphite(3) HSO3- Sulphite SO32- Hydrogen sulphate(4) HSO4- Sulphate SO42Phosphate PO43- (1) Also known as "bisulphide" or "disulphide". Also known as "bicarbonate". (3) Also known as "bisulphite". (4) Also known as "bisulphate". (2) Weak Chemical Bonds Play an Important Role in the Chemistry of Life Most chemical bonds in living organisms are covalent. Some chemicals in a cell contact forming a weak chemical bond. These weak bonds will respond to chemicals in another cell thus forming a response. This response is similar to brain cells responding to some stimulus. Hydrogen bonding occurs when a hydrogen atom bonds covalently to an electronegative atom that is also bonded to another electronegative atom. 12 The molecules which have this extra bonding are: Note: The solid line represents a bond in the plane of the screen or paper. Dotted bonds are going back into the screen or paper away from you, and wedge-shaped ones are coming out towards you. Notice that in each of these molecules: The hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge. Each of the elements to which the hydrogen is attached is not only significantly negative, but also has at least one "active" lone pair. Lone pairs at the 2-level have the electrons contained in a relatively small volume of space which therefore has a high density of negative charge. Lone pairs at higher levels are more diffuse and not so attractive to positive things. Another example of hydrogen bonding is water with ammonia. The electronegative atoms are nitrogen and hydrogen. The water orientated itself with the negative nitrogen atom forming a weak covalent bond. A non-polar molecule can have regions of “+” and “-“ causing weak bonds to be temporally formed. These weak attractions are Vander Waal forces. Biological Function Related to Shape The molecular shape of molecules (hybridization) determines they respond and recognize another molecule and react with it. Dalton’s Atomic Theory revolutionized chemistry by explaining chemical properties in terms of small, indivisible pieces of matter called atoms that are linked together to form polyatomic species (both ions and molecules). As chemists explored the properties of the polyatomic species, it became clear that they have size and 13 shape and that shape is particularly important in explaining their physical properties and why and how chemical reactions occur. A summary of some of the commonly observed polyatomic shapes (also known as structures or geometries) are shown in Table 1 arranged by the number of atoms around the "central atom." These are idealized structures; real molecules seldom exhibit these idealized shapes. However, the ideal shapes are good starting points toward understanding how the spatial arrangement of atoms in polyatomic species affect their properties and chemistry. Number of Atoms Around “Central Atom” Shape 1 Linear 2 Linear Bent Trigonal planar 3 Trigonal pyramidal 14 T-shape Tetrahedral 4 Square planar Bisphenoid (see-saw) 15 5 Trigonal bipyramidal Square pyramidal 6 Octahedral Chemical Reactions Make or Break Chemical Bonds Chemical reactions demonstrate the conservation of mass and energy. In any chemical reaction chemical bonds are broken, new substances are formed. When the products formed equal the breakdown of the reactants then chemical equilibrium happens. Now that we know the how and why of chemical bonding, we can look at some chemical reactions. Chemical reactions happen all around us: when we light a match, start a car, eat dinner or walk the dog. A chemical reaction is the pathway by which two 16 substances bond together. In fact we have already discussed several chemical reactions. One we have mentioned is the reaction of hydrogen with oxygen to form water. To write the chemical reaction you would place the reactants (the substances reacting) on the left with an arrow pointing to the the products (the substances being formed). Given this information, one might guess that the reaction to form water is written: H+O H2O However there are 2 problems with this chemical reaction. First, because atoms like to have full valence shells, single H or O atoms are rare (and unhappy) creatures. As we saw in the previous lesson, both hydrogen and oxygen react with themselves to form the molecules H2 and O2, respectively. These hydrogen and oxygen molecules are much more common. Given this correction, one might guess that the reaction looks like this: H2 + O2 H2O But we still have one problem. As written, this equation tells us that 1 hydrogen molecule (with 2 H atoms) reacts with 1 oxygen molecule (with 2 O atoms) to form 1 water molecule (with 2 H atoms and 1 O atom). In other words, we seem to have lost 1 O atom along the way! To write a chemical reaction correctly, the number of atoms on the left side of a chemical equation has to be precisely balanced with the atoms on the right side of the equation. How does this happen in our example? In actuality, the O atom that we 'lost' reacts with a 2nd molecule of hydrogen to form a second molecule of water. The reaction is therefore written: 2H2 + O2 2H2O In the chemical reaction above, the number in front of the molecule (called a coefficient) indicates how many molecules participate in the reaction. A simulation of the reaction can be viewed by clicking below (the atoms are represented as spheres in the animation: red = hydrogen, blue = oxygen): 17