Presentation is loading. Please wait.

Presentation is loading. Please wait.

Lecture 2 OrR.

Similar presentations

Presentation on theme: "Lecture 2 OrR."— Presentation transcript:

1 Lecture 2 OrR

2 Overview: Inquiring About the World of life
Evolution is the process of change that has transformed life on Earth. Biology is the scientific study of life Biologist ask questions such as: How does a single cell develop into an organism? How does the human brain work? How do living things interact in communities?

3 Evolution, the Overarching Theme of Biology
Evolution makes sense of everything we know about living organisms Organisms living on Earth are modified descendents of common ancestors

4 The Biological Hierarchy
The study of life can be divides into different level of biological organization

5 Life, Chemistry and Water
Biology is a multidisciplinary science Living organisms are subject to basic laws of physics and chemistry

6 Matter, Elements & compounds
Organisms are composed of matter Matter – anything that takes up space and has mass Mass measures the quantity of matter in an object and is defined by: Volume is how much space it takes up Density measure concentration of matter Matter is made up of elements An element is a substance that cannot be broken down to other substances by chemical reactions A compound is a substance consisting of two or more elements in a fixed ratio. A compound has characteristics different from those of its elements

7 Sodium Chloride(NaCl)
Table Salt Sodium Chloride(NaCl) Sodium (Na) Chlorine (Cl) Is oxygen an element or a compound? What about iron? Water? Carbon di oxide?

8 Essential Elements of Life
About 25 of the 92 elements are essential of life Carbon (C), hydrogen (H), oxygen (O), and nitrogen make up 96% of living matter. Most of the remaining consist of calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl) and magnesium (Mg). Trace elements are those required by an organism in minute quantities. Example: selenium, Zinc

9 Atomic Structure Each elements consists of one type of atom Each atom consists of an atomic nucleus surrounded by fast moving, negatively charged electrons Atomic nuclei contain positively charged protons – the number of protons (atomic number) identifies an element The nuclei of all atoms (except hydrogen) also contain uncharged neutrons

10 ISOTOPES OF CARBON Isotopes Makes atomic number not as simple as it may seem Protons and Neutrons in the nucleus Neutrons can vary independently of the number of protons Neutrons add weight to the atom Isotopes have the same number of protons, different number of neutrons

11 Isotopes are used in biological research as tracers
Isotopes in Research Radioisotopes decay can be used to estimate the age of organic material, rocks, or fossils that contain them. Isotopes are used in biological research as tracers Radioisotopes such as 14C, 32P, and 25S Stable, nonradioactive isotopes such as 15N (heavy nitrogen)

12 Radioactive iodine becomes concentrated in the thyroid gland
123I can be used to visualize the thyroid gland Scans of human thyroid glands after 123I was injected into the blood stream Radioactive iodine becomes concentrated in the thyroid gland

13 Chemical Bonds and Chemical Reactions
Atoms of reactive elements tend to combine into molecules by forming chemical bonds The four most important chemical linkages in biological molecules are ionic bonds, covalent bonds, hydrogen bonds and van der Waals forces. Chemical reactions occur when atoms or molecules interact to form new chemical bonds or break old ones

14 Chemical Bonds and Chemical Reactions
Forces: Covalent bonds Hydrogen bonds Ionic bonds van der Waals bonds Macromolecules Carbohydrate Proteins Nucleic acids Lipids

15 Covalent Bonds Principle form to hold atoms together
based on sharing electrons very strong

16 to break a C-C bond require 83 kcal/mole
the bond cant be break with any physiological condition The bond can only be bend, stretch or rotate Enzymes: Biological catalysts that enable specific bonds to be broken or formed under physiological conditions

17 Types of covalent bonds:
rotate Single bonds C-C Double bonds C=C Triple bonds C=C saturated Stronger unsaturated Can’t rotate Other molecules: Oxygen molecule O=O Nitrogen molecule N=N Nitrogen is strong molecule, cant be broken by any organism, but by bacteria which produce enzyme name nitrogenase

18 Chirality characteristics of covalent bond
Carbon can make 4 COVALENT bonds comes out as tetrahedron

19 Molecules are known as Optical isomer
Chirality characteristics of covalent bond Mirror image Molecules are known as Optical isomer Important concept: At molecular level much of biology relies on interaction of complementary 3d surfaces

20 Sharing of electron C-C or C-H equal sharing (non-polar bond) N-H, O-H unequal sharing (polar bond) Electronegativity

21 Two Types of Covalent Bonds
Nonpolar Covalent Bond The atoms participating in the bond are sharing electrons equally. There is no difference in charge between the two ends of such bonds. Polar Covalent Bond Atoms participating in the bonds do not share electrons equally. One atom pulls the electrons a little more toward its "end" of the bond, so that atom bears a slightly negative charge. The atom at the other end of the bond bears a slightly positive charge. Nonpolar Covalent Bond Examples: The bonds in molecular hydrogen (H2), oxygen (02), and nitrogen (N2) mentioned earlier are examples. These molecules are some of the gases that make up air. The bonds in methane (CH3), another gas, are also nonpolar. Polar Covalent bond Examples: For example, a water molecule (H—O—H,) has two covalent bonds; both are polar. The oxygen atom in a water molecule carries a slight negative charge, and each of the hydrogen atoms carries a slight positive charge. Any separation of charge into distinct positive and negative regions is called polarity.

22 Polar Covalent Bond Polarity occurs when atoms electrons unequally due to differences in electronegativities. This is seen in water (H2O). More electronegative atoms tend to pull electrons toward them creating a polar molecule.

23 The polarity of water molecules results in hydrogen bonding

24 Ionic Bonding Sodium chloride (table salt) is an example of ionic bonding, that is, electron transfer among atoms or redox reaction.

25 Ionization Molecules formed by ionic bonding breakup (ionization) when dissolved in water (solvent), producing separate positive (cation) and negative (anion) ions. These ions conduct electricity and thus called electrolytes.

26 Hydrogen Bonds A hydrogen bond is a weak attraction between a hydrogen atom and another atom taking part in a separate polar covalent bond Like ionic bonds, hydrogen bonds form by the mutual attraction of opposite charges. Unlike ionic bonds, hydrogen bonds do not make molecules out of atoms, so they are not chemical bonds Hydrogen bonds form and break much more easily Example: Hydrogen bonds form between water molecules. The hydrogen atom has a slight positive charge and the other atom has a slight negative charge. Hydrogen bonds form and break much more easily than covalent or ionic bonds do. Even so, many of them form. Their collective strength imparts unique properties to many substances such as water. Hydrogen bonds that form among the atoms of biological molecules such as DNA hold these molecules in their characteristic shapes.

27 Hydrogen Bond Hydrogen bonding is formed between the partially positive (hydrogen) end of a polar molecule and the negative end of another (e.g. O2 or N2). Example : Water molecules

28 Summary: From Atoms to Molecules

29 Water: The Molecule That Supports All of Life
Water is the biological solvent on Earth All living organisms require water more than any other substance Most cells are surrounded by water, and cells themselves are about 70-95% water The abundance of water is the main reason the Earth is habitable.

30 Water Water: Polar covalent bonds gives water its unique properties that make life possible Water is an excellent solvent Hydrophilic substances Hydrophobic substances Water is stable in high temperature Why ice floats on water Water has cohesive properties Water evaporates Water's special properties as a liquid begin with the two polar covalent bonds in each water molecule. Overall, the molecule has no charge, but the oxygen pulls the shared electrons a bit more than the hydrogen atoms do. Thus, each of the atoms in a water molecule carries a slight charge: The oxygen atom is slightly negative, and the hydrogen atoms are slightly positive. The separation of charge means that the water molecule itself is polar. The polarity is very attractive to other water molecules, and hydrogen bonds form between them in tremendous numbers. Extensive hydrogen bonding between water molecules imparts unique properties to liquid water, and those properties make life possible. First, water is an excellent solvent. A solvent is a liquid that can dissolve other substances. When a substance dissolves, its individual molecules or ions become solutes as they disperse. Salts, sugars, and many other compounds that dissolve easily in water are polar, so many hydrogen bonds form between them and water molecules. A salt is a compound that dissolves easily in water and releases ions other than H+ and OH" when it does. Hydrogen bonding with water dissolves such hydrophilic (water-loving) substances by pulling their individual molecules away from one another and keeping them apart. You can see how water interacts with hydrophobic (water-dreading) substances if you shake a bottle filled with water and salad oil, then set it on a table and watch what happens. Salad oil consists of nonpolar molecules, and water molecules do not form many hydrogen bonds with nonpolar molecules. Shaking breaks some of the hydrogen bonds that keep water molecules together. However, the water quickly begins to cluster into drops as new hydrogen bonds form among its molecules. The bonding excludes molecules of oil and pushes them together into droplets that rise to the surface of the water. The same interaction occurs at the thin, oily membrane that separates the water inside of cells from the water outside of them. The organization of membranes—and of life—starts with such interactions. A second property of water is temperature stability. Temperature is a way to measure the energy of molecular motion: All molecules jiggle nonstop, and they jiggle faster as they absorb heat. However, extensive hydrogen bonding restricts the movement of water molecules—it keeps them from jiggling as much as they would otherwise. Thus, compared with other liquids, water absorbs much more heat before its temperature rises. Temperature stability is an important component of homeostasis, because most of the molecules of life function properly only within a certain range of temperature. Below 0°C (32°F), water molecules do not jiggle enough to break hydrogen bonds, and they become locked in the rigid, lattice-like bonding pattern of ice. Individual water molecules pack less densely in ice than they do in water, so ice floats on water. During cold winters, ice sheets may form near the surface of ponds, lakes, and streams. Such ice "blankets" insulate liquid water under them, so they help keep fish and other aquatic organisms from freezing. A third life-sustaining property of liquid water is cohesion, which means that water molecules resist separating from one another. This property is important in many processes that sustain multicelled bodies. As one example, water molecules constantly escape from the surface of liquid water as vapor, a process called evaporation. Evaporation is resisted by the hydrogen bonding that keeps water molecules together. In other words, overcoming water's cohesion takes energy. Thus, evaporation sucks energy in the form of heat from liquid water, which decreases its surface temperature. Evaporative water loss can help you and some other mammals cool off when you sweat in hot, dry weather. Sweat, which is about 99 percent water, cools the skin as it evaporates. Take-Home Message: Why is water essential to life? Being polar, water molecules hydrogen-bond to one another and to other polar (hydrophilic) substances, and repel nonpolar (hydrophobic) substances. Extensive hydrogen bonding between water molecules gives water unique properties that make life possible: cohesion, temperature stability, and a capacity to dissolve many substances..

31 Diffusion

32 Diffusion

33 OSMOSIS Osmosis is the net diffusion of water down its own concentration gradient Water molecules can readily permeate the plasma membrane Slower Process Aquaporins One molecule of solute can displace one molecule of water

34 OSMOSIS Osmosis when pure water is separated from a solution containing a nonpenetrating solute. The net diffusion of water down its concentration gradient through a selectively permeable membrane is known as osmosis

Solutes that can penetrate the plasma membrane do not contribute to osmotic differences between the ICF and ECF and do not affect cell volume (although before equilibrium is achieved, transient changes in volume may occur as a result of differing rates of diffusion of water and the solute across the membrane

36 Tonicity The effect of a solution on the osmotic movement of H20. The tonicity of a solution has no units and is a reflection of its concentration of nonpenetrating solutes relative to the cell’s concentration of nonpenetrating solutes Isotonic Equal tension to plasma. RBCs will not gain or lose H20. Hypotonic: Osmotically active solutes in a lower osmolality and osmotic pressure than plasma. RBC will hemolyse. Hypertonic: Osmotically active solutes in a higher osmolality and osmotic pressure than plasma. RBC will crenate.


38 Acids and Bases What is pH
Acids: Fossil fuel, N2 containing fertilizers, Acid rain Bases Buffers Hydrogen ions contribute to pH. Acids release hydrogen ions in water; bases accept them. Salts release ions other than H+ and OH-. Buffers keep the pH of body fluids stable. They are part of homeostasis.

Download ppt "Lecture 2 OrR."

Similar presentations

Ads by Google