BIO 110- Chapter 2.pdf

General Biology Bio 110 Chapter 2 CONTENTS: Part 1 Life on Earth: An Overview Chapter 1 PART II Chemistry of Life a. Basic Chemistry Chapter 2 b. Chemistry of Organic Molecules Chapter 3 PART III The C...

General Biology Bio 110 Chapter 2 CONTENTS: Part 1 Life on Earth: An Overview Chapter 1 PART II Chemistry of Life a. Basic Chemistry Chapter 2 b. Chemistry of Organic Molecules Chapter 3 PART III The Cell a. Cell Structure and Function Chapter 4 b. Membrane Structure and Function Chapter 5  Part II CHEMISTRY OF LIFE a. Basic Chemistry 2.1 Chemical Elements a. Elements Matter only exists in three distinct states: solid liquid gas Matter is composed of certain basic substances called elements. Element is a substance that cannot be broken down to simpler substances by ordinary chemical means. There are 92 elements serving as the building blocks of matter; of which six (acronym CHNOPS) are basic to life and make up 95% of the body weight of organisms. They are hydrogen, oxygen, sulfur Carbon, nitrogen, and phosphorus. Other important elements include iron, potassium, magnesium and calcium. 2.1 Chemical Elements b. Atoms The atomic theory: atom is the smallest part of an element that displays the property of the element. Atomic symbol is made of one or two letters, i.e., H for hydrogen atom and Na for sodium atom. subatomic particles include: positively charged protons uncharged neutrons negatively charged electrons Protons and neutrons are located within the nucleus of an atom, while electrons move around the nucleus (ex. Helium, Figure 2.1). Figure 2.1. Model of Helium atom( He)  2.1 Chemical Elements c. Atomic Number and Mass Number: Atomic number is the number of protons housed in the atom nucleus. Protons and neutrons are assigned one atomic mass unit (AMU) each. Electrons are so small that their AMU is considered zero in most calculation. Therefore, the mass number of an atom is the sum of protons and neutrons in the nucleus. The term mass is used, and not weight, because mass is constant, while weight changes according to the gravitational force of a body. The gravitational force of the Earth is greater than that of the moon; therefore, substances weigh less on the moon, even though their mass has not changed. The atomic number is usually written as a subscript to the lower left of the atomic symbol (2He). The mass number is written as a superscript to the upper left of the atomic symbol (4He). Regardless of position, the smaller number is always the atomic number, as shown here for carbon.  d. The Periodic Table: The periodic table was constructed according to certain chemical and physical characteristics. The atoms shown in the periodic table (Figure 2.2) are assumed to be electrically neutral. The atomic number indicates the number of protons and number of electrons. To determine the number of neutrons, subtract the number of protons from the atomic mass, and take the closest whole number. Figure 2.2 A portion of the periodic table.  d. The Periodic Table (continued): Every atom is in a particular period (the horizontal rows) and in a particular group (the vertical columns). The atomic number of every atom in a period increases by one if you read from left to right. All the atoms in a group share the same binding characteristics. For example, all the atoms in group VII react with one atom at a time. The atoms in group VIII are called the noble gases because they are inert and rarely react with another atom. Figure 2.2 A portion of the periodic table. e. Isotopes: Isotopes are atoms of the same element that differ in the number of neutrons. Isotopes have the same number of protons, but they have different atomic masses. For example, the element carbon has three common isotopes: carbon 12 has six neutrons carbon 13 has seven neutrons, carbon 14 has eight neutrons. Unlike the other two isotopes of carbon, carbon 14 is unstable and can be changed over time into nitrogen 14, which is a stable isotope of the element nitrogen. As carbon 14 decays, it releases various types of energy in the form of rays and, therefore, it is a radioactive isotope.  Part IIa 2.2 COMPOUNDS AND MOLECULES 2.2 Compounds and Molecules A compound exists when two or more elements have bonded together, while molecule is the smallest part of a compound. These two terms are used interchangeably, but in biology, we usually use the term “Molecules”. Water (H2O) is a molecule that contains atoms of hydrogen and oxygen. A formula tells the number of each kind of atom in a molecule as in glucose. 2.2 Compounds and Molecules a. Ionic Bonding: Ionic bond also called electrovalent bond; type of linkage formed from electrostatic attraction between oppositely charged ions in a chemical compound.. Table salt( sodium chloride or NaCl ) is an example.. Sodium is a metal and chlorine is a poisonous gas.. However, when chemically combined, an edible combine emerges.  The Chemical Bonds: a. Ionic Bonding: Sodium (Na) has only one electron in its third shell and tends to be an electron donor (Figure 2.3). Once it gives up this electron, the second shell, with eight electrons, becomes its outer shell. Chlorine (Cl), on the other hand, tends to be an electron acceptor. Its outer shell has seven electrons. Figure 2.3 Formation of sodium So, if it acquires only one more electron, it chloride (table salt) will have a completed outer shell. Now both atoms have eight electrons in their outer shells.  a. Ionic Bonding (continued): This electron transfer causes a charge imbalance in each atom: The sodium atom has one more proton than it has electrons; therefore, it has a net charge of +1 (symbolized by Na+). The chlorine atom has one more electron than it has protons; therefore, it has a net charge of -1 (symbolized by Cl-). Such charged particles are called ions.  Ionic compounds are held together by an attraction between charged ions called ionic bonding. When such ions reacts, an ionic compound called sodium chloride (NaCl) results. Figure 2.3 Formation of sodium chloride (table salt) b. Covalent Bonding :  It results when two atoms share electrons in such a way that each atom has an octet (eight) of electrons in the outer shell. Ex., one hydrogen atom can share with another hydrogen atom (H—H for structural formula or H2 for molecular formula). Atoms can also share more than one pair of electrons to complete their octets. A double covalent bond occurs when two atoms share two pairs of electrons (Figure 2.4).  To show that oxygen gas (O2) contains a double bond, the molecule can be written as O=O. Figure 2.4 Covalently bonded molecules.  c. Polar and Non-polar Covalent Bonds: When the sharing of electrons between two atoms is equal, the covalent bond is said to be a non-polar covalent bond. If one atom can attract electrons to a greater degree than the other atom, it is the more electronegative atom. Electronegativity is dependent on the number of protons—the greater the number of protons, the greater the electronegativity. When electrons are not shared equally, the covalent bond is a polar covalent bond. Ex., the oxygen atom in water is more electronegative than the hydrogen atoms and the bonds are polar. https://saylordotorg.github.io/text_general-chemistry-principles-patterns- and-applications-v1.0/s12-09-polar-covalent-bonds.html  Part IIa 2.3 CHEMISTRY OF WATER  2.3 CHEMISTRY OF WATER The structural formula of water molecule (Figure 2.5, on the far left) shows that when water forms, an oxygen atom is sharing electrons with two hydrogen atoms. The ball-and-stick model (in the center) shows that the covalent bonds between oxygen and each of the hydrogens are at an angle of 104.5°. The space-filling molecule (on the far right) gives us the three-dimensional shape of the molecule and Figure 2.5 Water molecule. indicates its polarity. The shape of water, and all organic molecules, is necessary to the structural and functional roles they play in living things like a key fits a lock. The shape and polarity of a water molecule makes hydrogen bonding possible. a. Hydrogen Bonding The dotted lines in Figure 2.5 indicate that the hydrogen atoms in one water molecule are attracted to the oxygen atoms in other water molecules. This attraction is weaker than an ionic or covalent bond. The dotted lines indicate that hydrogen bonds are more easily broken than covalent bonds. Figure 2.5 Water molecule.  b. Properties of Water: Water represents 70–90% of all living cells. Without hydrogen bonding, water would boil at -91°C and melt at -100°C, making most of the water on Earth steam and life unlikely. Because of hydrogen bonding, water is a liquid at temperatures typically found on the Earth’s surface. It melts at 0°C and boils at 100°C. When examining the other planets with the hope to find life, we look for signs of water. 1. Water Has a High Heat Capacity and High Heat of Evaporation: Hydrogen bonding helps water absorb heat without a great change in temperature. Converting 1 g of the coldest liquid water to ice requires the loss of 80 calories of heat energy (Figure 2.6). Figure 2.6 Temperature and water.  b. Properties of Water Water holds onto its heat and its temperature falls more slowly than that of other liquids. This property of water is important not only for aquatic organisms but also for all living things. This is because hydrogen bonds must be broken before water boils. This property gives animals in a hot environment an efficient way to release excess body heat. When an animal sweats, body heat is used to vaporize water, thus cooling the animal (Figure 2.6). Figure 2.6 Temperature and water.  b. Properties of Water 2. Water is a universal solvent: Due to its polarity, water facilitates chemical reactions, both outside and within living systems. It dissolves a great number of substances. When ionic salts - for example, sodium chloride (NaCl) - are put into water, the negative ends of the water molecules are attracted to the sodium ions and the positive ends of the water molecules are attracted to the chloride ions. This causes the sodium ions and the chloride ions to separate, or dissociate, in water. Molecules that can attract water are said to be hydrophilic. When ions and molecules disperse in water, they allow reactions to occur. Nonionized and non-polar molecules that cannot attract water are said to be hydrophobic, ex., gasoline, and, therefore, it does not mix with water.  b. Properties of Water: 3- Cohesive and Adhesive Nature. ( leading to Capillary Action) Water cohesion refers to the ability of water molecules to cling to each other due to hydrogen bonding. Because of cohesion, water exists as a liquid under ordinary conditions of temperature and pressure. Water adhesion refers to the ability of water molecules to cling to other polar surfaces. This is because of water polarity. Animals often contain internal vessels in which water assists the transport of nutrients and wastes because the cohesion and adhesion of water allows blood to fill the tubular vessels of the cardiovascular system. Cohesion and adhesion also contribute to the transport of water in plants. Water evaporating from the leaves is immediately replaced with water molecules from transport vessels that extend from the roots to the leaves. Accordingly, water tension is created to pull the water column up from the roots. Adhesion of water to the walls of the vessels also helps prevent the water column from breaking apart..  b. Properties of Water: A combination of adhesive and cohesive forces accounts for capillary action, tendency of water to move in narrow tubes, even against the force of gravity (Figure 2.7). In a narrow tube (Figure 2.7a), there is adhesion between the water molecules and the glass wall of the tube. Capillary action occur when the adhesion is stronger than cohesive forces between the liquid and molecules. Other water molecules inside the tube are then “pulled along” because of cohesion, which is due to hydrogen bonds between the water molecules. In the wider tube (Figure 2.7b), a smaller percentage of the water molecules line the glass wall. As a result, the adhesion is not strong enough to overcome the cohesion of the water molecules beneath the surface level of the container and water in the tube rises only slightly. Water moves through the spaces between soil particles to the roots of plants by capillary action. Figure 2.7 Capillary action.  Part II a 2.4 ACIDS AND BASES  2.4 ACIDS AND BASES When water ionizes, it releases an equal number of hydrogen ions (H+) (also called a proton) and hydroxide ions (OH-): a. Acidic Solutions (High H+ concentration): Lemon juice, vinegar, tomatoes, and coffee are acidic solutions. Acids are substances that dissociate in water, releasing hydrogen ions (H+). The acidity of a substance depends on how fully it dissociates in water. Ex., HCl, a strong acid, dissociates almost completely in this manner: HCl → H+ + Cl- If HCl is added to water, number of Figure 2.8 The pH scale. hydrogen ions (H+) increases.  b. Basic Solutions (Low H+ Concentration): Magnesia and ammonia are basic solutions. Bases are substances that either take up hydrogen ions (H+) or release hydroxide ions (OH-). Ex., NaOH, a strong base, dissociates in this manner: NaOH → Na+ + OH- If NaOH is added to water, number of hydroxide ions (OH-) increases. c. pH Scale: The pH scale is used to indicate the acidity or basicity (alkalinity) of a solution. The pH scale in Figure 2.8 ranges from 0 to 14. Figure 2.8 The pH scale.  c. pH Scale (continued): A pH of 7 represents a neutral state in which the hydrogen (H+) ion and hydroxide (OH-) ion concentrations are equal. A pH below 7 is an acidic solution because H+ concentration is greater than OH- concentration. A pH above 7 is basic because OH- concentration is greater than H+. As we move down the pH scale from pH 14 to pH 0, each unit is 10 times more acidic than the previous unit. As we move up the scale from 0 to 14, each unit is 10 times more basic than the previous unit. Therefore pH 5 is 100 times more acidic than is pH 7 and 100 times more basic than pH 3. Figure 2.8 The pH scale. Thank You