Biology Textbook PDF

Biology Sylvia S. Mader Michael Windelspecht...

Biology Sylvia S. Mader Michael Windelspecht Dr. Qamar Abbas Solangi Copyright © McGraw-Hill Education. Permission required for reproduction or display. 1  Basic Chemistry Copyright © McGraw-Hill Education. Permission required for reproduction or display. 2  Outline Introduction 2.1 Chemical Elements 2.2 Molecules and Compounds 2.3 Chemistry of Water 2.4 Acids and Bases 3  What is Biology Biology is the scientific study of life. There is great diversity among living things. Living things  are composed of the same chemical elements as nonliving things.  obey the same physical and chemical laws that govern everything in the universe. 4  Characteristics of Life Living things are organized. ( from atom to biosphere) Life requires materials and energy ( Sun: Source of energy, metabolism: all chemical reactions in cell) Living things maintain homeostasis (Homeostasis is the maintenance of internal conditions) Living things respond to stimuli. Living things reproduce and develop. Living things have adaptations. (that may lead to evolution) 5  2.1 Chemical Elements Matter refers to anything that has mass and occupies space. Matter exists in four states: solid, liquid, gas and plasma. All matter (both living and non-living) is composed of basic substances called elements. 6  Elements All matter (living or non-living) is composed of basic substance called elements. An element is a substance that cannot be broken down into simpler substances by ordinary chemical means; composed of one type of atom. Each element has its own unique property like density, solubility, melting point, and reactivity Ninety-two elements are naturally occurring that serve as building blocks Six elements make up 95% of the body weight of organisms (acronym CHNOPS):  Carbon  Hydrogen  Nitrogen  Oxygen  Phosphorus 7  Sulfur Composition of Earth’s Crust and Its Organisms 8  Atoms An atom is the smallest part of an element that displays the property of the element (Gk, “indivisible”) John Dalton’s atomic theory says that elements consists of tiny particles called atoms An element and its atom share the same name. Each element is represented by one or two letters to give it a unique atomic symbol.  H = hydrogen, Na = sodium, C = carbon 9  Atoms  Composed of subatomic particles: protons, neutrons, electrons  Central nucleus Protons – positively charged, 1 amu Neutrons – no charge, 1 amu  Orbiting clouds around nucleus (electron shells) Electrons – negatively charged, very low mass-negligible in calculations 10 Atoms He 11 Atoms-2 He 12 Atoms-3 He 13 Jump to long image description Atomic Number and Mass Number The atomic number is equal to the number of protons in each atom of an element. The mass number of an atom is equal to the sum of the number of protons and neutrons in atom’s nucleus. It is an average mass for all isotopes of that atom e.g. C has 12.01 Term mass is used not weight as mass is constant while weight changes according to gravitational force of body  The atomic mass is approximately equal to the mass number. 14 Atomic Number and Mass Number 15  Isotopes Isotopes are atoms of the same element that differ in the number of neutrons (and therefore different atomic masses). Some isotopes of an element are unstable or radioactive  Some isotopes spontaneously decay. Carbon 14 is an example of a radioactive isotope. – Has been used to examine reactions in photosynthesis Carbon 12 Carbon 13 Carbon 14 16  Periodic Table Dmitri Mendeleev Atoms of an element are arranged horizontally by increasing atomic number in rows called periods. Atoms of an element arranged in vertical columns are called groups. Atomic number in a period increases by 1 from L to R.  Atoms within the same group share the same chemical binding characteristics.  Group VIII are the noble gases and are inert. Atoms shown in the periodic table are electrically neutral.  Therefore, the atomic number tells you the number of electrons as well as the number of protons. 17 All the atoms in a particular group have the same number of valence electrons and therefore share common chemical characteristics. Each period shows the number of electron shells 18 for an element.  Electrons and Energy Electrons are attracted to the positively charged nucleus; thus, it takes energy to hold electrons in place. It takes energy to push them away and keep them in their own shell.  The more distant the shell, the more energy it takes to hold in place. Electrons have energy due to their relative position (potential energy). Electrons determine chemical behavior of atoms. 19  The Distribution of Electrons The Bohr model is a useful way to visualize electron location.  Electrons revolve around the nucleus in energy shells (energy levels).  For atoms with atomic numbers of 20 or less, the following rules apply: The first energy shell can hold up to 2 electrons. Each additional shell can hold up to 8 electrons. Each lower shell is filled first before electrons are placed in the next shell.  These rules cover most of the biologically significant elements. 20  The Distribution of Electrons Periods tell that how many shells an atom have Groups tell that how many electrons in outer shell (valence shell)  Electrons revolve around the nucleus in energy shells (energy levels).  For atoms with atomic numbers of 20 or less, the following rules apply: The first energy shell can hold up to 2 electrons. Each additional shell can hold up to 8 electrons. Each lower shell is filled first before electrons are placed in the next shell.  These rules cover most of the biologically significant elements. 21  Valence Electrons The outermost energy shell of any atom is called the valence shell. The valence shell is important because it determines many of an atom’s chemical properties. The octet rule states that the outermost shell is most stable when it has eight electrons.  Exception: If an atom has only one shell, the outermost valence shell is complete when it has two electrons. 22  Valence Electrons-1 The number of electrons in an atom’s valence shell determines whether the atom gives up, accepts, or shares electrons to acquire eight electrons in the outer shell. –Atoms that have their valence shells filled with electrons tend to be chemically stable. –Atoms that do not have their valence shells filled with electrons are chemically reactive. 23 Bohr Models of Atoms Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Jump to long image description 24  2.2 Molecules and Compounds A molecule is two or more elements bonded together.  It is the smallest part of a compound that retains its chemical properties. NaCl, H , etc. A compound is a molecule containing at least two different elements bonded together.  CO , H O, C H O , etc. A formula tells the number of each kind of atom in a molecule.  C H O means there is one molecule of glucose 25 containing 6C,12H, 6O. Compounds and Molecules 26  Chemical Bonding Bonds that exist between atoms in molecules contain energy. Organisms are directly dependent on chemical-bond energy Bonds between atoms are caused by the interactions between electrons in outermost energy shells. The process of bond formation is called a chemical reaction. 27 Types of Bonds: Ionic Bonding An ion is an atom that has lost or gained an electron. e.g Na+, K+, Mg+2, Ca+2, Cl-1 An ionic bond forms when electrons are transferred from one atom to another atom and the oppositely charged ions are attracted to each other.  Example: formation of sodium chloride Salts are solid substances that usually separate and exist as individual ions in water. 28 Formation of Sodium Chloride 29 Formation of Sodium Chloride 30 Formation of Sodium Chloride 31 Formation of Sodium Chloride-3. In a sodium chloride crystal, ionic bonding between Na+ and Cl− causes the atoms to assume a three-dimensional lattice in which each sodium ion is surrounded by six chloride ions, and each chloride ion is surrounded by six 32 sodium ions. The result is crystals of salt as in table salt. Types of Bonds: Covalent Bonds Covalent bonds result when two atoms share electrons so each atom has an octet of electrons in the outer shell.  Note: In the case of hydrogen, the outer energy shell is complete when it contains two electrons. 33 Covalently Bonded Molecules 34 Types of Bonds: Covalent Bonds In a nonpolar covalent bond electrons are shared equally between atoms.  Examples: hydrogen gas, oxygen gas, methane  Methane is nonpolar due to its symmetry In a polar covalent bond electrons are shared unequally.  Example: water, -NH2 (amine group) Electronegativity is the ability of an atom to attract electrons towards itself in a chemical bond.  In water, the oxygen atom is more electronegative than 35 the hydrogen atoms and the bonds are therefore polar. Types of Bonds: Covalent Bonds 36 2.3 Chemistry of Water 37  2.3 Chemistry of Water Water is a polar molecule.  The shape of a water molecule and its polarity make hydrogen bonding possible. This is an example of a structure-function relationship. A hydrogen bond is a weak attraction between a slightly positive hydrogen atom and a slightly negative atom.  It can occur between atoms of different molecules or within the same molecule.  A single hydrogen bond is easily broken while multiple hydrogen bonds are collectively quite strong.  It helps to maintain the proper structure and function of 38 complex molecules such as proteins and DNA.  Hydrogen bonding  It can occur between atoms of different molecules or within the same molecule.  A single hydrogen bond is easily broken while multiple hydrogen bonds are collectively quite strong.  It helps to maintain the proper structure and function of complex molecules such as proteins and DNA.  H bond holds two strands of DNA H bond is weaker than ionic or covalent bonds Dotted lines show that these bonds are easily broken 39  Properties of Water Water molecules cling together because of hydrogen bonding.  This association gives water many of its unique chemical properties.  Without H bonding----water could freeze at -100C and boils at -91C Water has a high heat capacity.  The presence of many hydrogen bonds allow water to absorb a large amount of thermal energy without a great change in temperature. 40  Water has a high heat capacity  A calorie is the amount of heat energy needed to raise the temperature of 1 g of water 1ͦC  Converting 1 g of the coldest liquid water to ice require the loss of 80 calories of heat energy  Water holds on its heat and its temperature falls more slowly than any other liquid Allows organisms to maintain their normal internal temperatures, protected from rapid temperature changes. 41 Temperature and Water 42  Properties of Water When water boils, it evaporates Converting 1 g of the hottest water to a gas requires 540 calories Water has a high heat of evaporization.  Hydrogen bonds must be broken to evaporate water.  Bodies of organisms cool when their heat is used to evaporate water. Example: sweating 43  Water is a good solvent  Water is a good solvent because of its polarity.  Polar substances dissolve readily in water.  A solution contains dissolved substances, or solutes.  Hydrophilic molecules dissolve in water  Hydrophobic molecules do not dissolve in water e.g. nonpolar molecules like Gasoline  It dissolves anionic salt Sodium Chloride (NaCl) and Ammonia (NH3) 44 Water as a Solvent 45  Properties of Water-2 Water molecules are cohesive and adhesive.  Cohesion is the ability of water molecules to cling to each other due to hydrogen bonding. Water flows freely Water exists as a liquid under temperature and pressure of earth. Surface tension , stronger the force between molecules. Greater the surface tension Water strider (insect) can walk on water 46  Properties of Water-2 Water molecules are cohesive and adhesive.  Adhesion is the ability of water molecules to cling to other polar surfaces. Due to water’s polarity Capillary action  Cohesion and adhesion account for water transport in plants as well as transport in blood vessels. 47 Water as a Transport Medium Water evaporates, pulling the water column from the roots to the leaves. Water molecules cling together and adhere to sides of vessels in stems. Water enters a plant at root cells. 48  Properties of Water-3 Frozen water (ice) is less dense than liquid water.  At temperatures below 4°C, hydrogen bonds between water molecules become more rigid but also more open.  Water expands as it reaches 0°C and freezes.  Ice floats on liquid water. Without this property, ice would sink and oceans freeze solid, instead of from the top down. Acts as an insulator on top of a frozen body of water. 49 Ice is less dense than water 50 Jump to long image description A Pond in Winter 51  2.4 Acids and Bases pH is a measure of hydrogen ion concentration in a solution. When water ionizes or dissociates, it releases an equal number of hydrogen (H ) ions and hydroxide (OH ) ions. Acids are substances that dissociate in water, releasing hydrogen ions. e.g. lemon juice, vinegar, tomatoes, coffee HCl H+ + Cl- 52  2.4 Acids and Bases Bases are substances that either take up hydrogen ions (H ) or release hydroxide ions (OH ). NaOH Na+ + OH- Milk of magnesi (MgOH) and ammonia 53  The pH Scale The pH scale is used to indicate the acidity or basicity (alkalinity) of a solution.  Values range from 0–14 0 to 7 to 14 = Basic (or alkaline)  Logarithmic scale Each unit change in pH represents a 10-fold change in H concentration pH of 4 is 10X as acidic as pH of 5 pH of 10 is 100X more basic than pH of 8 54 The pH Scale-1 Jump to long image description 55  Buffers and pH A buffer is a chemical or a combination of chemicals that keeps pH within normal limits. Health of organisms requires maintaining the pH of body fluids within narrow limits.  Human blood is normally pH 7.4 (slightly basic). If blood pH drops below 7.0, acidosis results. If blood pH rises above 7.8, alkalosis results. Both are life-threatening situations. 56  Buffers and pH  Body has built-in mechanisms to prevent pH changes. Example: carbonic acid buffer dissociates and re-forms to reduce changes in pH H2CO3 H+ + HCO-3 H+ + HCO-3 H2CO3 OH- + H2CO3 HCO-3 + H2O 57