ORGANIC CHEMISTRY. Serkan SAYINER, DVM PhD, Assist. Prof.

Transcription

1 ORGANIC CHEMISTRY Serkan SAYINER, DVM PhD, Assist. Prof.

2 ENERGY AND MATTER The Units of Energy, Energy and Nutrition, The Three States of Matter, Classification of Matter, Intermolecular Forces, Boiling and Melting Point, Energy and Phase Changes, Gases Serkan SAYINER, DVM PhD, Assist. Prof.

3 ENERGY Energy is the capacity to do work. Potential energy is stored energy. Kinetic energy is the energy of motion. Energy can be converted from one form to another, one rule, the law of conservation of energy, governs the process. The total energy in a system does not change. Energy cannot be created or destroyed.

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5 THE UNITS OF ENERGY Energy can be measured using two different units, calories (cal) and joules (J). A calorie is the amount of energy needed to raise the temperature of 1 g of water 1 C. Joules and calories are related. 1 cal = J Since both the calorie and the joule are small units of measurement, more often energies in reactions are reported with kilocalories (kcal) and kilojoules (kj). 1 kcal = 1,000 cal 1 kj = 1,000 J 1 kcal = kj

6 ENERGY and NUTRITION When an organism eats food, The protein, carbohydrates, and fat (lipid) in the food are metabolized to form small molecules, That in turn are used to prepare new molecules that cells need for maintenance and growth. This process also generates the energy needed for the organs to function, allowing the heart to beat, the lungs to breathe, and the brain to think.

7 ENERGY and NUTRITION The amount of stored energy in food is measured using nutritional Calories (upper case C), Where 1 Cal = 1,000 cal. Since 1,000 cal = 1 kcal, the following relationships exist. Upon metabolism, proteins, carbohydrates, and fat each release a predictable amount of energy, the caloric value of the substance. For example, one gram of protein or one gram of carbohydrate typically releases about 4 Cal/g, while fat releases 9 Cal/g. If we know the amount of each of these substances contained in a food product, we can make a first approximation of the number of Calories it contains by using caloric values as conversion factors.

8 ENERGY and NUTRITION When an individual eats more Calories than are needed for normal bodily maintenance, the body stores the excess as fat. The average body fat content for men and women is about 20% and 25%, respectively. This stored fat can fill the body s energy needs for two or three months. Frequent ingestion of a large excess of Calories results in a great deal of stored fat, causing an individual to be overweight.

9 THE THREE STATES OF MATTER Matter exists in three common states - gas, liquid, and solid. A gas consists of particles that are far apart and move rapidly and independently from each other. A gas has no definite shape or volume. The particles of a gas expand to fill the volume and assume the shape of whatever container they are put in.

10 THE THREE STATES OF MATTER A liquid consists of particles that are much closer together but are still somewhat disorganized since they can move about. The particles in a liquid are close enough that they exert a force of attraction on each other. A liquid has a definite volume, but takes on the shape of the container it occupies.

11 THE THREE STATES OF MATTER A solid consists of particles - atoms, molecules, or ions - that are close to each other and are often highly organized. The particles in a solid have little freedom of motion and are held together by attractive forces. A solid has a definite volume, and maintains its shape regardless of the container in which it is placed.

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14 THE THREE STATES OF MATTER Whether a substance exists as a gas, liquid, or solid depends on the balance between the kinetic energy of its particles and the strength of the interactions between the particles. In a gas, the kinetic energy of motion is high and the particles are far apart from each other. As a result, the attractive forces between the molecules are negligible and gas molecules move freely. In a liquid, attractive forces hold the molecules much more closely together, so the distance between molecules and the kinetic energy is much less than the gas. In a solid, the attractive forces between molecules are even stronger, so the distance between individual particles is small and there is little freedom of motion.

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16 CLASSIFICATION of MATTER All matter can be classified as either a pure substance or a mixture. A pure substance is composed of a single component and has a constant composition, regardless of the sample size and the origin of the sample. A pure substance cannot be broken down to other pure substances by any physical change.

17 CLASSIFICATION of MATTER A mixture is composed of more than one component. The composition of a mixture can vary depending on the sample. A mixture can be separated into its components by physical changes. Dissolving table sugar in water forms a mixture, whose sweetness depends on the amount of sugar added. If the water is allowed to evaporate from the mixture, pure table sugar and pure water are obtained. Mixtures can be formed from solids, liquids, and gases.

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19 CLASSIFICATION of MATTER A pure substance is classified as either an element or a compound. An element is a pure substance that cannot be broken down into simpler substances by a chemical reaction. A compound is a pure substance formed by chemically combining (joining together) two or more elements. Nitrogen gas, aluminum foil, and copper wire are all elements. Water is a compound because it is composed of the elements hydrogen and oxygen.

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22 INTERMOLECULAR FORCES Ionic compounds are composed of extensive arrays of oppositely charged ions that are held together by strong electrostatic interactions. These ionic interactions are much stronger than the forces between covalent molecules, so it takes a great deal of energy to separate ions from each. In covalent compounds, the nature and strength of the attraction between individual molecules depend on the identity of the atoms.

23 INTERMOLECULAR FORCES Intermolecular forces are the attractive forces that exist between molecules. There are three different types of intermolecular forces in covalent molecules, presented in order of increasing strength: 1.London dispersion forces (van der Waals forces) 2.Dipole dipole interactions 3.Hydrogen bonding The strength of the intermolecular forces determines whether a compound has a high or low melting point and boiling point.

24 INTERMOLECULAR FORCES London dispersion forces are very weak interactions due to the momentary changes in electron density in a molecule. All covalent compounds exhibit London dispersion forces. The larger the molecule, the larger the attractive force between two molecules, and the stronger the intermolecular forces. Dipole dipole interactions are the attractive forces between the permanent dipoles of two polar molecules. Hydrogen bonding occurs when a hydrogen atom bonded to O, N, or F, is electrostatically attracted to an O, N, or F atom in another molecule.

25 BOILING POINT and MELTING POINT The boiling point (bp) of a compound is the temperature at which a liquid is converted to the gas phase, while the melting point (mp) is the temperature at which a solid is converted to the liquid phase. The strength of the intermolecular forces determines the boiling point and melting point of compounds. The stronger the intermolecular forces, the higher the boiling point and melting point.

26 BOILING POINT and MELTING POINT Methane (CH 4 ) and water (H 2 O) are both small molecules with hydrogen atoms bonded to a second-row element, so you might expect them to have similar melting points and boiling points. Methane, however, is a nonpolar molecule that exhibits only London dispersion forces, whereas water is a polar molecule that can form intermolecular hydrogen bonds. As a result, the melting point and boiling point of water are much higher than those of methane. In fact, the hydrogen bonds in water are so strong that it is a liquid at room temperature, whereas methane is a gas. In comparing two compounds with similar types of intermolecular forces, the larger compound generally has more surface area and therefore a larger force of attraction, giving it the higher boiling point and melting point.

27 BOILING POINT and MELTING POINT

28 ENERGY and PHASE CHANGES When energy is absorbed, a process is said to be endothermic. When energy is released, a process is said to be exothermic. In a phase change, the physical state of a substance is altered without changing its composition.

29 ENERGY and PHASE CHANGES Converting a Solid to a Liquid Converting a solid to a liquid is called melting. Melting is a phase change because the highly organized water molecules in the solid phase become more disorganized in the liquid phase, but the chemical bonds do not change. Each water molecule is composed of two O-H bonds in both the solid and the liquid phases. Melting is an endothermic process. Energy must be absorbed to overcome some of the attractive intermolecular forces that hold the organized solid molecules together to form the more random liquid phase. The amount of energy needed to melt 1 g of a substance is called its heat of fusion (cal/g).

30 ENERGY and PHASE CHANGES Freezing is the opposite of melting; that is, freezing converts a liquid to a solid. Freezing is an exothermic process because energy is released as the faster moving liquid molecules form an organized solid in which particles have little freedom of motion.

31 ENERGY and PHASE CHANGES Converting a Liquid to a Gas Converting a liquid to a gas is called vaporization. Vaporization is an endothermic process. Energy must be absorbed to overcome the attractive intermolecular forces of the liquid phase to form gas molecules. The amount of energy needed to vaporize 1 g of a substance is called its heat of vaporization (cal/g). Condensation is the opposite of vaporization; that is, condensation converts a gas to a liquid. Condensation is an exothermic process because energy is released as the faster moving gas molecules form the more organized liquid phase.

32 ENERGY and PHASE CHANGES A high heat of vaporization means that a substance absorbs a great deal of energy as it is converted from a liquid to a gas. Water has a high heat of vaporization. As a result, the evaporation of sweat from the skin is a very effective cooling mechanism for the body.

33 ENERGY and PHASE CHANGES Converting a Solid to a Gas Occasionally a solid phase forms a gas phase without passing through the liquid state. This process is called sublimation. The reverse process, conversion of a gas directly to a solid, is called deposition. Carbon dioxide is called dry ice because solid carbon dioxide (CO 2 ) sublimes to gaseous CO 2 without forming liquid CO 2. Carbon dioxide is a good example of a solid that undergoes this process at atmospheric pressure. At reduced pressure other substances sublime. For example, freeze-dried foods are prepared by subliming water from a food product at low pressure.

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35 GASES Anyone who has ridden a bike against the wind knows that even though we can t see the gas molecules of the air, we can feel them as we move through them. Air is a mixture of; 78% nitrogen (N 2 ), 21% oxygen (O 2 ), and 1% other gases, including carbon dioxide (CO 2 ), argon (Ar), water (H 2 O), and ozone (O 3 ).

36 GASES Simple gases in the atmosphere oxygen (O 2 ), carbon dioxide (CO 2 ), and ozone (O 3 ) are vital to life. Oxygen, which constitutes 21% of the earth s atmosphere, is needed for metabolic processes that convert carbohydrates to energy. Green plants use carbon dioxide, a minor component of the atmosphere, to store the energy of the sun in the bonds of carbohydrate molecules during photosynthesis. Ozone forms a protective shield in the upper atmosphere to filter out harmful radiation from the sun, thus keeping it from the surface of the earth.

37 GASES Properties of Gases A gas consists of particles atoms or molecules that move randomly and rapidly. The size of gas particles is small compared to the space between the particles. Because the space between gas particles is large, gas particles exert no attractive forces on each other. The kinetic energy of gas particles increases with increasing temperature. When gas particles collide with each other, they rebound and travel in new directions. When gas particles collide with the walls of a container, they exert a pressure.

38 GASES Pressure (P) is the force (F) exerted per unit area (A). Pressure = Force : Area = F : A All of the gases in the atmosphere collectively exert atmospheric pressure on the surface of the earth. The value of the atmospheric pressure varies with location, decreasing with increasing altitude. Atmospheric pressure also varies slightly from day to day, depending on the weather. Atmospheric pressure is measured with a barometer. A barometer consists of a column of mercury (Hg) sealed at one end and inverted in a dish of mercury.

39 GASES Many different units are used for pressure. The two most common units are the atmosphere (atm), and millimeters of mercury (mm Hg), where 1 atm = 760. mm Hg. One millimeter of mercury is also called one torr. In the United States, the common pressure unit is pounds per square inch (psi), where 1 atm = 14.7 psi. Pressure can also be measured in pascals (Pa), where 1 mm Hg = Pa.

40 GASES Taking a patient s blood pressure is an important part of most physical examinations. Blood pressure measures the pressure in an artery of the upper arm using a device called a sphygmomanometer. A blood pressure reading consists of two numbers such as 120/80, where both values represent pressures in mm Hg.

41 GASES The higher number is the systolic pressure and refers to the maximum pressure in the artery right after the heart contracts. The lower number is the diastolic pressure and represents the minimum pressure when the heart muscle relaxes. A desirable systolic pressure is in the range of mm Hg. A desirable diastolic pressure is in the range of mm Hg.

42 GASES Four variables are important in discussing the behavior of gases; pressure (P), volume (V), temperature (T), and number of moles (n). The relationship of these variables is described by equations called gas laws that explain and predict the behavior of all gases as conditions change. Three gas laws illustrate the interrelationship of pressure, volume, and temperature.

43 GASES 1. Boyle s law relates pressure and volume. 2. Charles s law relates volume and temperature. 3. Gay Lussac s law relates pressure and temperature.

44 GASES Boyle s law: For a fixed amount of gas at constant temperature, the pressure and volume of a gas are inversely related. p x V = Consant (k) p: Pressure V: Volume p 1 V 1 =p 2 V 2

45 GASES Charles s Law: All gases expand when they are heated and contract when they are cooled. For a fixed amount of gas at constant pressure, the volume of a gas is proportional to its Kelvin temperature. Volume and temperature are proportional; that is, as one quantity increases, the other increases as well. Thus, dividing volume by temperature is a constant (k) V : T = k Increasing the temperature increases the kinetic energy of the gas particles, and they move faster and spread out, thus occupying a larger volume. V 1 :T 1 =V 2 :T 2

46 GASES Gay Lussac s Law: Gay Lussac s law describes how the pressure of a gas changes as the Kelvin temperature is changed. For a fixed amount of gas at constant volume, the pressure of a gas is proportional to its Kelvin temperature. Pressure and temperature are proportional; that is, as one quantity increases, the other increases. Thus, dividing the pressure by the temperature is a constant (k). P:T=k Increasing the temperature increases the kinetic energy of the gas particles, and if the volume is kept constant, the pressure exerted by the particles increases. P 1 :T 1 =P 2 :T 2

47 GASES The Combined Gas Law: All three gas laws - Boyle s, Charles s, and Gay Lussac s laws - can be combined in a single equation, the combined gas law, that relates pressure, volume, and temperature. P 1 V 1 T 1 = P 2 V 2 T 2

48 GASES Avogadro s Law: Avogadro s law describes the relationship between the number of moles of a gas and its volume. When the pressure and temperature are held constant, the volume of a gas is proportional to the number of moles present. As the number of moles of a gas increases, its volume increases as well. Thus, dividing the volume by the number of moles is a constant (k). The value of k is the same regardless of the identity of the gas. V : n = k Thus, if the pressure and temperature of a system are held constant, increasing the number of moles increases the volume of a gas. V 1 :n 1 =V 2 :n 2

49 GASES The Ideal Gas Law: All four properties of gases - pressure, volume, temperature, and number of moles - can be combined into a single equation called the ideal gas law. The product of pressure and volume divided by the product of moles and Kelvin temperature is a constant, called the universal gas constant and symbolized by R. PV=nRT

50 GASES Dalton s Law: It describes the relationship between the partial pressures of the components and the total pressure of a gas mixture. The total pressure (P total ) of a gas mixture is the sum of the partial pressures of its component gases. Thus, if a mixture has three gases (A, B, and C) with partial pressures PA, PB, and PC, respectively, the total pressure of the system (P total ) is the sum of the three partial pressures. P total = P A + P B + P C

51 REFERENCE BOOKS Smith JG (2010). Organic Chemistry, 3rd Edition, McGraw-Hill. Smith JG (2012). General, Organic, & Biological Chemistry 2nd Edition, McGraw-Hill.