General Chemistry for Pharmaceutical Sciences Part I Final (2)

 General Chemistry for Pharmaceutical Sciences PHARM-101 Presented by Dr. Azza H. Rageh Associate Professor of Pharmaceutical Analytical Chemistry Office: College of Pharmacy Department of Pharmacognosy and Pharmaceutical Chemistry Building 15 – O...

 General Chemistry for Pharmaceutical Sciences PHARM-101 Presented by Dr. Azza H. Rageh Associate Professor of Pharmaceutical Analytical Chemistry Office: College of Pharmacy Department of Pharmacognosy and Pharmaceutical Chemistry Building 15 – Office 106 References  Raymond Chang (2015). Chemistry: 12th ed., McGraw- Hill education, Boston, New York, U.S.A.  Janice Smith (2022). General, Organic & biological Chemistry, 5th edition McGraw-Hill education, Boston, New York, U.S.A.  Theodore E. Brown, H. Eugene LeMay, Bruce E. Bursten, Catherine Murphy, Patrick Woodward, Matthew E. Stoltzfus (2019), Chemistry: The Central Science, 13th ed. , Pearson, UK. A- Matter and Measurements 1- Classification and States of Matter The Chemical World Chemistry: The science that seeks understanding the properties and behavior of matter by studying atoms and molecules. - Chemistry is central to understand many other scientific fields. - Virtually, everything around you is composed of “chemicals”. 5 Chemistry: The Central Science 6 Atoms and Molecules ⮚ Atoms are the building-blocks of matter. ⮚ Each element is made of a unique kind of atoms (so far, 120 elements are identified in the universe, they are represented in the periodic table of elements). ⮚ The compound Ñ is made of two or more atoms of different elements, bonded together to form molecules (molecules are the building-blocks of compounds). ⮚ The properties of a substance are determined by the properties of its constituent molecules and atoms. 7 Atoms and Molecules Important Note: some elements are present as “molecules” instead of “free atoms”, they are called: “Molecular Elements”, such as: H2, N2 , O2 , F2 , Cl2 , Br2 , I2 , P4 , S8 , Se8 and O3 (Ozone Gas) 8 Atoms and Molecules Example 1 ✓ The air contains carbon monoxide pollutant. ✓ Each molecule contains a carbon atom and an oxygen atom held together by a chemical bond. 9 Atoms and Molecules Example 2 Note: Balls of different colors are used to represent atoms of different elements. Attached balls represent connections between atoms that are seen in nature. These groups of atoms are called molecules. 10 The Classifications of Matter Matter is anything that occupies space and has mass. Examples: your textbook, your desk, the air around you, and even your body, are all composed of matter. Matter is everything around us. Matter can be classified according to: 1. State (the physical form) 2. Composition (the components that make it up) 11 The States of Matter ⮚ Matter can exist in one of three main states: solid, liquid, or gas. The state of matter changes from solid to liquid to gas by increasing temperature, and vice versa! 12 Solid Matter ⮚ Solid Matter: is composed of tightly packed particles (atoms or molecules). Solids retain their shapes because the particles are not free to move. ⮚ Although the atoms and molecules vibrate in solids, they do not move around or past each other. ⮚ Consequently, solid matter has a fixed (definite) volume and a fixed (rigid) shape. Examples of solids: Ice, aluminum, iron, wood, salt, and diamond. 13 Solid Matter Crystalline or Amorphous? ⮚ Crystalline Solids: atoms or molecules are arranged in “patterns” with a long-range repeating order. Important Examples on crystalline solids: table salt (NaCl) and diamond. ⮚ Amorphous Solids: atoms or molecules are not arranged in long-range patterns. Important Examples on amorphous solids: graphite, rubber, glass and plastic. 14 Liquid Matter ⮚ Liquid Matter: is made of more loosely packed particles than in solids. Particles can move about within a liquid, but they are packed densely enough that volume is maintained. ⮚ The ability of liquids to flow, makes them take the shapes of their containers. ⮚ Liquids have fixed volume but no fixed shape. Examples of liquids: water, oil, and gasoline. 15 Gaseous Matter ⮚ Gaseous Matter: is composed of particles packed so loosely that it has neither a defined shape nor a defined volume. ⮚ Particles of gases (atoms or molecules) are free to move relative to one another. ⮚ Gases have no fixed volume and no fixed shape, they take the volume and shape of their containers. These qualities make gases compressible. ⮚ Examples of gases: oxygen, nitrogen, CO2, water vapor 16 Summary of State Changes of Matter 17 Classification of Matter according to its composition ⮚ Matter can be divided into two classes: 1. Mixtures: are composed of more than one substance and can be physically separated into its component substances. 2. Pure substances: are composed of only one substance and can NOT be physically separated. 18 Mixtures There are two types of mixtures: 1. Heterogeneous mixtures 2. Homogeneous mixtures i.WS ✓ Heterogeneous Mixture: does NOT have uniform properties throughout. – (sand + water), (oil + water) or (gasoline + water) are examples on heterogeneous mixtures. ii A ✓ Homogeneous Mixture: has uniform properties throughout. – (salt water), (sugar + water) and alloys are homogeneous mixtures. 19 Pure Substances There are two types of pure substances: 1. Compounds 2. Elements ✓ Compound: can be chemically separated into individual elements. There are millions of compounds in the universe. ⮚ Water is a compound that can be separated into hydrogen and oxygen. ✓ Element: cannot be broken down further by chemical reactions. ⮚Elements are the 120 members of the periodic table of elements, such as: Sodium, Iron, Gold, Silver, Hydrogen, Oxygen, Carbon …... etc 20  Summary of Types of Matter Matter can be classified according to its composition into: pure substances (elements or compounds) and mixtures (homogeneous or heterogeneous): 21 Assessment 1) The process is which a solid substance is transformed directly into a gas is calledsublimationand it requires of temperature. increasing 2) Condensing is the physical process which changes a gas into a liquid, and it needs of temperature. decreasing 3) Which state of matter has a fixed volume but not a fixed shape? liquid 4) A matter is able to assume both the shape and volume of its container. gas 5) The ability of both gas and liquid states of matter to flow makes them able to change their shape to the shape of their reservoir. 6) Classify each substance as a pure substance or a mixture, and indicate the type of Pure substance min.tkgenous a. sweat rectifies b. carbon dioxide snItsgncek msYiatagge c. aluminum d. salt compound e. rust mixtures Puresubstance 1164 YHeferogenous f. wet sand g.Homogenous air Element h. oxygen gas i. bronze alloy j. honey 22 mittmogenous minthogenous  A- Matter and Measurements 2- Physical and Chemical Changes and properties 23 Physical and Chemical Changes Physical Changes: A process that does NOT cause a substance to become a different 64 substance (i.e. only the appearance (state or shape) is changed, but NOT the chemical composition). w̅ at Physical changes are reversible. Example 1: when water (H2O) boils, it changes its state from liquid to gas. ⮚ The gas remains composed of H2O, so this is a “physical change”. Example 2: when a piece of paper is shredded, or a glass window is broken, only their shapes have changed, but their chemical compositions remained unchanged. 24 Physical and Chemical Changes Chemical Changes: A process that causes a substance to change into a new substance with a new chemical composition. During a chemical change, atoms rearrange themselves to make different substances. GIA Chemical changes are irreversible. Example 1: rusting of iron is a chemical change: 4 Fe + 3 O2 → 2 Fe2O3 Example 2: burning of gasoline produces CO2 + H2O, so, it’s a chemical change 25 Evidences for Chemical Changes a) Release of a gas (e.g. bubbles or smoke) b) Emission of light or heat (e.g. burning of wood) c) Permanent change in color (e.g. the brown layer of iron rust) 26 Physical and Chemical Changes Examples 27 Physical and Chemical Properties of Matter Physical Properties: any characteristic that can be measured without changing the substance’s chemical identity or composition (i.e. without any chemical reactions). ⮚ Examples on Physical Properties: ⮚ Color ⮚ Viscosity ⮚ Odor ⮚ Temperature ⮚ Taste ⮚ Hardness ⮚ Density ⮚ Metallic Luster ⮚ Melting Point ⮚ Malleability ⮚ Boiling Point ⮚ Ductility 28 Physical and Chemical Properties of Matter Physical Properties: a r e e ith e r e xte n s ive o r in te n s i ve  Extensive properties: are that depend on amount of matter to be measured such as mass and volume.  Insensitive properties: are that not depend on amount of matter to be measured such as color, taste and density.  Classify the following physical properties extensive or intensive. 1. Color (intensive). 2. Density (intensive). 3. Volume (extensive). 4. Mass (extensive). 5. Boiling point (intensive): temperature at which substance boils. 6. Melting point (intensive): temperature at which substance melts. 29 Physical and Chemical Properties of Matter Chemical Properties: any characteristic that can be measured only by changing a substance’s chemical identity or composition (i.e. in a chemical reaction). ⮚ Examples on Chemical Properties: ⮚ Reactivity with other chemicals (acids, water, oxygen, ….) ⮚ Acidity and Basicity ⮚ Flammability: ability of compound to burn when exposed to flame. ⮚ Chemical stability: refers to reactions that alter chemical structures of compounds such as oxidation (reaction with oxygen), hydrolysis (reaction with water) and photosensitivity (decomposition by light). ⮚ Toxicity ⮚ Heat of combustion: Energy released upon complete combustion (burning) of compound with oxygen. ⮚ Oxidation state: It describes the degree of oxidation (loss of electrons) of 30 an atom in a chemical compound Assessment Identify the following as chemical or physical changes or properties: 1. blue colorPhysical Property 2. melting pointPhysicalProperty 3. density Physical P roperty chemical change 4. reaction with water 5. flammabilitychemicalProperty 6. hardness PhysicalProperty 7. toxicity chemicalProtenty 8. boiling point physicalproperty 9. reaction with acid chemicalchar 10. luster PhysicalProperty 11. perfume odor PhysicalProperty12. sour taste Physical Property 13. coal Burnschemicalchange 14. dry ice sublimes Physicalchange 15. Ag (Silver) tarnisheschemicalchange 16. milk sourschemicalchange 17. an apple is cut Physicalchange 18. fruit rot chemicalchange 19. heat changes H2O to steam Physicalcharge 20. pancakes cook chemical change 21. baking soda reacts to vinegarchemicalchange22. grass grows chemicalchange 23. iron rustschemicalchange 24. a tire is inflated Physicalchange 25. alcohol evaporates Physicalchange 26. food is digested chemicalchange 27. ice melts Physicalchange 28. paper absorbs water Physicalchange 31  A- Matter and Measurements 3- Units of Measurements and Density of Materials 32 The Units of Measurement We use measurements in everyday life, for example: walking 2.25 km to the university campus, carrying a backpack with a mass of 12 kg, and observing when the outside temperature has reached 40°C. 33 The Units of Measurement w̅ ⮚ Units: standard quantities used to specify measurements, they a r e critical in chemistry. The most common systems of units are: 1. The English system: used in the United States. 2. The Metric system: used in most countries. 3. The International System of Units (SI): used by scientists, and it is based on the metric system. 34 Units in the Metric and SI Systems of Measurement ⮚ In the metric and SI systems, one unit is used for each type of measurement: Measurement Metric System (SI) System Length meter (m) meter (m) Volume liter (L) cubic meter (m3) Mass gram (g) kilogram (kg) Temperature Celsius (°C) Kelvin (K) Time second (s) second (s) 35 The Meter: A Measure of Length (or Distance) Length (or distance): ⮚ Useful relationships between ▪ is measured using a the units of length: meter stick. 1m 391in ▪ The unit meter (m) is used in both the metric and SI systems. ▪ Centimeters (cm) is used for smaller lengths. 36 The Kilogram: A Measure of Mass The mass of an object is a measure of the quantity of matter within it. The SI unit of mass is kilogram (kg): 1 kg = 2.21 lb (pound) 1 gram = 1/1000 kg = (10-3 kg) Weight of an object is a measure of the gravitational pull on its matter: (weight ≠ mass) 37 Units for Volume Measurement ⮚ The common units for volume measurements are: Liter (L), Milliliter (mL), and Cubic Meter (m3) ⮚ Useful relationships between the units of volume:  1 L = 1000 mL  1000 L = 1 m3 38 Some Lab Tools for Volume Measurement ⮚ Volume is the amount of space occupied by a substance. 39 Unit of Time Measurement Time measurement: ▪ uses the unit second (s) in both the metric and SI systems. Days, Hours, Minutes, Seconds ⮚ Useful relationships between the units of time: 40 Unit of Temperature Measurement ⮚ Common Units of Temperature: – Fahrenheit (oF) (English system). – Celsius (oC) (Metric system). – Kelvin (K) (SI system). ⮚ Example 1: Boiling Point of Water: 212 oF = 100 oC = 373.15 K ⮚ Example 2: Freezing Point of Water: 32 oF = 0 oC = 273.15 K 0 K is called: “absolute zero”. It is the lowest possible temperature in the universe! 41 Relationships Between Units of Temperature K C 273 15 - From °C to K: K = °C + 273.15 - From K to °C: R 273.15 °C = K – 273.15 - From °C to °F: F C 18 C 32 °F = [1.8 × (°C)] + 32 - From °F to °C: F FX8 37 c 13 5 42 Examples on Temperature Conversions 1. How much does 350 oF equal in both oC and K? 32 3,5 177C oC = (350 − 32)/1.8 = 318/1.8 = 177 oC R 1774 233.15 450K K = 177 + 273.15 = 450.15 K 2. Convert ( ̶ 40 oC) to oF: F 40 1 8 32 405 oF = [1.8 × (−40)] + 32 = −72 + 32= −40 oF 688 31 3. Express 298 Kelvin in degree Celsius: C 298 3273.15 24.856 oC = 298 − 273.15 = 24.85 oC 43 Prefix Multipliers: Changing The Value of The Unit ⮚ The International System of Units (SI) uses the prefix multipliers with the standard units. These prefix multipliers change the values of the units (make units larger or smaller). Examples: - the kilometer has the prefix kilo, meaning 1000 meter (or 103 m). - the millimeter has the prefix milli, meaning 1/1000 meter (or 10−3 m). 44 Prefix Multipliers: Increasing the size of the Unit Prefixes that INCREASE the size of unit: G a Note: students shall memorize those prefixes: (names, symbols, and factors)! 45 Prefix Multipliers: Decreasing the size of the Unit Prefixes that DECREASE the size of unit: Note: students shall memorize those prefixes: (names, symbols, and factors)! 46 Density of Materials ▪ Material's Density is its mass per unit volume. ▪ Usually measured in g/L, g/mL, or g/cm3. ⮚ Density Expression: Useful Note: 1 mL = 1 cm3 47 Calculating Density – Example If a 0.258 g sample of HDL has a volume of 0.215 cm3, what is the density (in g/cm3), of the HDL? 0.258 0.215 1.291cm Step 1: State the given and needed quantities: Given Needed 0.258 g HDL density of HDL in g/cm3 0.215 cm3 HDL Step 2: Use the relation: 48 Calculating Density – Example 1. Do the following conversions: 55 103 0.055 103 1 a. 55 m = ……………0.055 km 5500 cm = …………… 11 103 b. 11 s = …………… 11000 ms 1100=1 …………… 11 153 0all ks 2.7 c. 2.7 g = 162.7 1,12 pg …………… = 182.7 ……………153 ng d. 3.6 L = …………… 3600 mL = 3.6 19 µL …………… 1 C 56,2321 48.8 2. Express the temperature −56 °F in both °C and K. K 48.88 273.15224.27K 3. Perform each of the following unit conversions: 13.53Mt p 14.74yd 2.87kg fig 6.342lb a. 13.53 m to yd b. 2.87 kg to lb FId.Fi3xiixii 2 4524 123tcmxfgum4.87in c. 2.45 L to m3 123.7 mm to in 4. Calculate the density of penny that has a mass of 2.49 g and a volume of 0.349 cm3. D 7 7.1391cm 49  B- Atoms, Molecules, Ions and Periodicity 1- History of Atomic Structure, Modern Atomic Theory and Atomic structure  History of Atomic Structure ▪ The word atom comes from the ancient Greek adjective atomos, meaning "indivisible”. indivisible : unable to ▪ Atom is smallest constituent of matter and be divided or separated consists of nucleus and surrounding electrons. ▪ The nucleus is made of one or more protons and typically a similar number of neutrons. ▪ Every atom is composed of a nucleus and one or more electrons bound to the nucleus 1 4h I 51 History of Atomic Structure ▪ Protons and neutrons are called nucleons. More than 99.94% of an atom's mass is in the nucleus. ▪ The protons have a positive electric charge, the electrons have a negative electric charge, and the neutrons have no electric charge. ▪ If the number of protons and electrons are equal, that atom is electrically neutral. ▪ If an atom has more or fewer electrons than protons, then it has an overall negative or positive charge, respectively, and it is called an ion. 52 History of Atomic Structure 53 How Electrons are Arranged in Atom? Electrons are arranged in orbitals (s, p, d, and f). An orbital can hold a maximum of two electrons. Orbitals are grouped in shells designated by numbers 1, 2, 3, and so on. Valence electrons are located in the outermost shell. Electrons in shells that are not completely filled. 54 How Electrons are Arranged in Atom? Electron arrangement (orbital diagram and electron configuration) of: 1H Is 10Ne 2He 15 15P 6C 16S 7N 17Cl 8O 18Ar Example: Orbital Diagram Electron configuration and orbital diagram of fluorine 9F is 252ps Electron configuration it A TL 9F 55 Modern Atomic Theory and Laws that Led to It  The theory that all matter is composed of atoms grew out of many observations and laws.  The three most important laws that led to the development and acceptance of the atomic theory are: Law of conservation of mass To Law of definite proportions IN.mn v6 Law of multiple proportions wait got 56 Law of Conservation of Mass Law of Conservation of Mass (A. Lavoisier):  Matter is neither created nor destroyed in a chemical reaction. Total mass of used reactants = Total mass of produced products Total number of reactants’ atoms = Total number of products’ atoms 57 Law of Definite Proportions Law of Definite Proportions (J. Proust):  A given chemical compound always contains the same elements in the exact same proportions (by mass), regardless to its source or how it was prepared For example: Sodium chloride (NaCl) always has a definite mass-to- mass ratio of chlorine and sodium. This ratio is always the same for any sample of pure NaCl, regardless of its origin: A 100 g sample of NaCl contains 39.3 g Na & 60.7 g Cl 𝑴𝒂𝒔𝒔 𝑪𝒍 𝟔𝟎.𝟕 𝒈 = = 𝟏. 𝟓𝟒 𝑴𝒂𝒔𝒔 𝑵𝒂 𝟑𝟗.𝟑 𝒈 A 58.44 g sample of NaCl contains 22.99 g Na & 35.44 g Cl 𝑴𝒂𝒔𝒔 𝑪𝒍 𝟑𝟓.𝟒𝟒 𝒈 = = 𝟏. 𝟓𝟒 𝑴𝒂𝒔𝒔 𝑵𝒂 𝟐𝟐.𝟗𝟗 58 Law of Multiple Proportions Law of Multiple Proportions (J. Dalton):  When two elements, “A” and “B”, combine with each other to form two or more compounds, the ratios of the masses of those elements in the formed compounds are simple whole numbers.  For example: – A molecule of carbon dioxide (CO2) has a ratio of 1 C atom to every 2 atoms of oxygen, or 1:2. – A molecule of carbon monoxide (CO) has a ratio of 1 C atom to 1 atom of oxygen, or 1:1.  Another Example: “Fe” to “O” in FeO = (1:1), while in Fe2O3 = (2:3) 59 Dalton’s Atomic Theory Postulates of Atomic Theory of Matter (J. Dalton):  Each element is composed of tiny, indestructible particles called atoms.  An element’s atoms are identical in size, mass and all other properties.  Atoms of different elements vary in size and mass.  Atoms of one element cannot change into atoms of another element.  Atoms are indivisible and the chemical reaction leads to rearrangement of atoms not to their creation or destruction.  Molecules are simple whole-number ratios of the combined elements 60 Elements: Defined by their Number of Protons Each element has a unique name, symbol, and atomic number. – Symbol has either one or two letters: O (oxygen) or Fe (iron) – The elements are arranged on the periodic table in order of increasing their atomic numbers. 61  Elements: Defined by their Number of Protons The number of protons located in an atom’s nucleus determines the element’s identity. – The number of protons in the nucleus of an atom is called the “atomic number” and is referred to as “Z”, it’s considered as the “fingerprint” of any element. 62 Isotopes: When The Number of Neutrons Varies I in Isotopes: are atoms of one element that have the same number of protons (atomic number) and different number of neutrons. – Isotopes differ in mass number because they have different number of neutrons. – Isotopes are chemically identical, but may be physically different. Mass Number (A) = Protons + Neutrons Note: Isotopes are identified by their “mass numbers” (example: C–12 , C–13 , C–14) 63 Isotopes: An Example Carbon-12 Carbon-13 Carbon-14 12C 13C 14C 6 6 6 protons: 6 p+ 6 p+ 6 p+ neutrons: 6 n 7n 8n electrons: 6 e- 6 e- 6 e- 64 Exercise How many protons, electrons, and neutrons are in the following atoms: protons electrons neutrons 32 S 6 16 16 16 65 Cu 29 36 29 29 U–240 Note: Neutral atoms have the same number of electrons as protons! 65 Exercise The Answer: How many protons, electrons, and neutrons are in the following atoms: protons electrons neutrons 32 S 16 p 16 e 32  16 = 16 n 16 65 Cu 29 p 29 e 65  29 = 36 n 29 92 U–240 92 p 92 e 240  92 = 148 n (Isotope) 66  Molecules ▪ Molecules are neutrally charged species and has no charge e.g. HCl if compared to ions. ▪ A molecule is an electrically neutral group of two or more atoms held together by chemical bonds. Molecules are distinguished from ions by their lack of electrical charge. ▪ A molecule may be homonuclear, that is, it consists of atoms of one chemical element as with oxygen (O2); or it may be heteronuclear, a chemical compound composed of more than one element, as with water (H2O). 67 Ions: Charged Atoms (Cations vs. Anions) Cations Anions A CATION forms when an atom loses An ANION forms when an atom gains one or more electrons from its outer one or more electrons into its outer shell (valence shell, the highest shell (valence shell, the highest energy level). energy level). Cations are positively charged Anions are negatively charged because the atom has more protons because the atom has fewer protons (+) than electrons (–). (+) than electrons (–). – Mg atom has 12 protons & 12 electrons. – F atom has 9 protons & 9 electrons. – Mg2+ ion has 12 protons & 10 electrons. – F – ion has 9 protons & 10 electrons. Metal elements tend to form cations. Nonmetal elements tend to form anions. Example: Mg  Mg2+ + 2 e – Example: F + 1 e–  F– 68 Ions: Charged Atoms (Cations vs. Anions) metal non metal 69 Ions: Zwitter Ion (Charged Molecules) Zwitter ions: Neutral molecules although it has positive and negative ions at different positions Ex. 1: most amino acids O O +NH -CH -COO- 3 2 Ex. 2: Pregabalin A medication used for treatment of epilepsy and anxiety 70 Assessment Answer the following questions: 1- Fill in the blanks to complete the table: 2- Determine the number of p+, n0, and e¯ in each atom: F7 10 7 D Pell e 11 n 12 1 186 e 86 n136 F 82,0 82 in 126 3- Determine the number of protons and the number of electrons in each ion: F23 e 4- Write isotopic symbols of the form for each isotope: a. the copper isotope with 36 neutrons b. the oxygen isotope with 8 neutrons 16 c. the aluminum isotope with 14 neutrons d. the iodine isotope with 74 neutrons 8 A I 71 B- Atoms, Molecules, Ions and Periodicity 2- The Periodic Table of Elements The Modern Periodic Table In 1913, Henry Moseley proposed the modern periodic table using atomic number instead of atomic mass, as the organizing principle for all the identified elements. The Modern Periodic Table Consists of:  7 Rows: are referred to as Periods, the periods are numbered 1–7.  18 Columns: are sometimes referred to as Groups or Families, they are numbered 1–18 (or the A and B grouping). – They are commonly called “Families” because the element within the column have similar physical and chemical properties. 73 Classification of Elements Elements in the periodic table are classified into the following three major divisions:  Metals  Nonmetals  Metalloids 74 The Modern Periodic Table: Metals, Nonmetals & Metalloids 75 Classification of Elements: Metals  Metals lie on the lower left side and middle of the periodic table.  Properties of Metals:  They are good conductors of heat and electricity.  All metals are solids at room temperature, except mercury (Hg) is a liquid.  They can be pounded into flat sheets (malleability).  They can be drawn into wires (ductility).  They are often shiny.  They tend to lose electrons when they undergo chemical changes (forming cations). About 75% of the elements in the periodic table are metals. 76 Classification of Elements: Nonmetals  Nonmetals lie on the upper right side of the periodic table.  Properties of Nonmetals:  Poor conductors of heat and electricity.  Can be found in all three states of matter (gases, liquids & solids).  Nonmetals with solid state are brittle (not ductile & not malleable).  They tend to gain electrons when they undergo chemical changes (forming anions). 77 Classification of Elements: Metalloids  Metalloids are elements that lie along the zigzag line that divides metals and nonmetals in the periodic table.  Properties of Metalloids:  Can exhibit mixed properties of both metals and nonmetals.  Solids at room temperature.  Known as semiconductors for electricity.  Poor conductors of heat. 78 Major Families: Alkali Metals (Group 1A) Group 1A elements are called the alkali metals, they are highly reactive metals (except hydrogen). Examples: – A marble-sized piece of sodium explodes violently when dropped into water. – Lithium, potassium, and rubidium are also alkali metals. 79 Major Families: Alkaline Earth Metals (Group 2A) is Group 2A elements are called the alkaline earth metals. They are fairly reactive, but not quite as reactive as the alkali metals (group 1A). Examples: – Calcium, for example, reacts fairly vigorously with water. – Other alkaline earth metals include magnesium (a common low-density structural metal), strontium, and barium. 80 Major Families: Halogens (Group 7A) Group 7A elements are called halogens, they are very reactive nonmetals. They are always found in nature as salts. Examples: – Chlorine, a greenish-yellow gas with a pungent odor. – Bromine, a red-brown liquid that easily evaporates into a gas. – Iodine, a purple solid. – Fluorine, a pale-yellow gas. 81 Major Families: Noble Gases (Group 8A) Group 8A elements, called the noble gases, are mostly unreactive (inert). Examples: – The most familiar noble gas is helium, used to fill buoyant balloons. – Other noble gases are neon (often used in electronic signs), argon (a small component of our atmosphere), krypton, and xenon. 82 Ions and the Periodic Table Loss electrons Gain electrons +1 +2 +3 -3 -2 -1 In general, the charge of ions of main-group elements can be predicted from their group’s number 83 Assessment Answer the following questions: 1- Which pairs of elements do you expect to be similar? Why? a. N and Ne b. Mo and Sr c. Ar and Kr d. Cl and I e. P and Pd 2- Predict the charge of the monoatomic ion formed by each element: a. O 2 b. K A c. Al 3 d. Rb 1 e. N 3 3- Using a copy of the periodic table, write the name of each element and classify it as a metal, nonmetal, or metalloid: a. Na b. Mg c. Br d. N e. As Sudiom metal magnesium nonmetal nonmetal metalloids metal 7- Using a copy of the periodic table, classify each element as an alkali metal, alkaline earth metal, halogen, or noble gas: a. sodium b. iodine c. calcium d. barium e. krypton alkali alkaliearth A nobel h halogene 84 gas  C- Molecules, Compounds and Chemical Bonds 1- Chemical Formulas and Molecular Models Elements, Compounds and Mixtures 86 Elements & Compounds Am  Elements are the simplest form of matter, they can combine together to make a limitless number of compounds. I  The properties of the compounds are totally different from their constituent elements.  Example: the properties of water are different from the properties of both H2 and O2: 87 Classification of Organic Compounds According to Functional groups Rg lb 0 CE 0H C O C 88 Classification of Organic Compounds According to Functional groups ohh 5 9 L c white oh Mt Of ENE 89 Representing Compounds: Chemical Formulas & Molecular Models A Compound is a distinct substance that is composed of bonded atoms of two or more elements. We can describe the compound by describing the number and type of each atom in the simplest unit of the compound: molecules Each element is represented by its letter symbol (from the periodic table) The number of atoms of each element is written to the right of the element as a subscript (if there is only one atom of an element, the “1” subscript is not written), examples: C6H12O6 , CH3Br Polyatomic ions are placed in parentheses (if more than one). examples: Mg(NO3)2 , Al(OH)3 , NaOH 90 Representing Compounds: Chemical Formulas & Molecular Models ItW Compounds are generally represented with their Chemical Formulas or Molecular Models. Chemical formula indicates the type and number of each element present in the compound (using the letter symbols of the elements from the periodic table): – Water is represented as H2O – Carbon dioxide is represented as CO2 – Sodium chloride is represented as NaCl – Carbon tetrachloride is represented as CCl4 91 Types of Chemical Formulas  Chemical formulas can generally be categorized into three different types: 1. Empirical Formula 2. Molecular Formula 3. Structural Formula The amount of information about the structure of the compound varies with the type of formula. All formulas and models convey a limited amount of information – none is a perfect representation! 92 Types of Chemical Formulas: The Empirical Formula mp.it  Empirical Formula: gives the relative number of atoms of each element in a compound. 6H1206 H2O  It does not describe the actual number of atoms, the order of attachment, or the shape of molecules.  It is the simplest whole-number ratio representation of the type and number of elements present in a molecule.  The Ionic compounds (metal + nonmetal) are usually represented using their Empirical Formulas (formula units): (e.g. MgO not Mg2O2 & CaS not Ca2S2 , …. ) 93 Types of Chemical Formulas: The Molecular Formula exact  Molecular Formula: gives the actual number of atoms of each see element in a molecule of a compound.  It does not describe the order of attachment, or the shape of molecules.  Examples: 4m16s a) H2O is the molecular formula of water, which means that the water molecule is actually composed of 2 hydrogen atoms + 1 oxygen atom. b) C4H8 is a molecular formula, which means that the molecule is actually empirical fcomposed of 4 carbon atoms + 8 hydrogen atoms. Hz c) B2H6 is a molecular formula, which means that the molecule is actually BH3 composed of 2 boron atoms + 6 hydrogen atoms. 94 Types of Chemical Formulas: The Structural Formula  Structural Formula: is a sketch or diagram of how the atoms in the molecule are bonded to each other. 5261161  It uses lines to represent covalent bonds and shows how atoms in a molecule are connected or bonded to each other.  lines describe the number of electrons shared by the bonded atoms: Single line = two shared electrons, a single covalent bond Double line = four shared electrons, a double covalent bond Triple line = six shared electrons, a triple covalent bond  It’s used only with “Molecular Compounds” (Nonmetal + Nonmetal), but NOT with “Ionic Compounds” (Metal + Nonmetal) aol.gs Ef  Example: The structural formula for CO2 is:  Example: The structural formula for methane CH4 is: 95 Molecular Models There are two common molecular models used to represent the molecules of compounds: we - Ball-and-Stick Model EI - Space-Filling Model Example: the different ways to represent the methane molecule (CH4): 96 Exercises: Molecular and Empirical Formulas assessment CsHr L Exercise: Write the Empirical Formulas for the following compounds: CHU CH CH CzHsCl CH Br 97 An Atomic Level View of Elements and Compounds 1. Atomic Elements: elements whose particles are single atoms, most of the elements in the periodic table are “atomic elements”: e.g. Fe , Na , Al , Ne , Hg, ….. 2. Molecular Elements: elements whose particles are multi-atom molecules, having the same type of atoms (i.e. atoms of the same element), (atoms are bonded by covalent bond): e.g. H2 , O2 , N2 , Cl2 , P4 , S8 , Se8 …. (see the next slide) 3. Molecular Compounds: compounds whose particles are molecules made of only nonmetals, (bonded by covalent bond): e.g. H2O , NH3 , HCl , CH4 4. Ionic Compounds: compounds whose particles are composed of cations (of metals) and anions (of nonmetals), (bonded by ionic bond): metai nonmetal nonmetal metal e.g. NaCl , AlF3 , Fe2O3 , Mg2S metal Eye nonmdetal fetal 98 Molecular Elements Certain elements occur as diatomic molecules (only 7 out of the 118 elements of the periodic table):  H2 , N2 , O2 , F2 , Cl2 , Br2 and I2 Some other elements occur as polyatomic molecules:  P4 , S8 , Se8 and O3 (Ozone Gas) 99 Classification of Elements and Compounds: A Summary 100 Assessment 6 Classify the following substances as: Atomic Elements, Molecular Elements, Molecular Compounds, or Ionic Compounds: a. Barium, Ba Atomicelement ………………………………………………..……………………………..……. b. Iron(III) chloride, FeCl3 …………………..………………..……………………… Inoic compound c. Bromine, Br2 …………………..…………………………………………………….………… Molecular element d. Ethanol, C2H6O …………………..………………..…………………………..……...…… Molecularcompound e. Nitrogen monoxide, NO molecular compound …………………..………………………….……..……… f. Cobalt, Co Atomicelement …………………..………………..………………………..……..……...…………… g. Carbon monoxide, CO …………………..……………………….……….………..… molecularcompound h. Nickel(II) chloride, NiCl2, Ioniccompound …………………..……………………..……………..… i. Sodium iodide, NaI …………………..………………………………….……………..… Ioniccompound j. Phosphorus chloride, PCl3 molecularcompound …………………..……………………..……………..… 101  B- Molecules, Compounds and case Chemical Bonds 2- Mole, Molar Mass and Mole conversions 102 Molar Mass  Definition of the “Mole”  The Molar Mass of a Compound  Mole Conversions 103 Formula Mass & Molar Mass AINT Formula Mass (amu): The mass of an individual molecule or formula unit, expressed in “amu” (atomic mass unit)  also known as molecular I mass or molecular weight.  Sum of the masses of the atoms in a single molecule or formula unit Formula mass of H2O = [2 × (1.01 amu H)] + [1 × (16.00 amu O)] = 18.02 amu Molar Mass (g/mol): The mass of one mole of a substance, expressed in “g/mol”  Molar mass is numerically equal to formula mass, but expressed in g/mol Molar mass of H2O = 18.02 g/mol 104 How to Calculate the Molar Mass of Any Substance? Ñ  To calculate the molar mass of any substance, you must first D know its exact chemical formula!  The molar mass can be calculated for any substance by summation of the atomic masses (from the periodic table) of all the atoms of elements present this substance’s formula: An element’s molar mass in grams per mole (g/mol) is numerically Is equal to the element’s atomic mass in atomic mass units (amu). a jim in  Examples: Calculate the molar mass (g/mol) for: - Molar mass of water (H2O) = (2×1) + (1×16) = 18 g/mol we - Molar mass of oxygen (O2) = 2×16 = 32 g/mol - Molar mass of NaCl = (1×23) + (1×35) = 58 g/mol - Molar mass of glucose (C6H12O6) = (12×6) + (12×1) + (16×6) = 180 g/mol 105  B- Molecules, Compounds and can Chemical Bonds 3- Chemical Reactions, Chemical equations, Intramolecular Forces, Types of Chemical Bonds, 106 Chemical Reactions A chemical reaction is a process that leads to the chemical transformation of one set of chemical substances to another.  The chemical reaction can be described by a chemical equation.  The substance (or substances) initially involved in a chemical reaction are called reactants or reagents.  The chemical reaction yields one or more products, which usually have properties different from the reactants. A + B AB Reactants Product 107 Chemical Equations  Chemical Equation: is a shorthand way of describing a chemical reaction. Provides some basic information about the reaction: – formulas of reactants and products – states of reactants and products – relative numbers of reactant and product molecules that are required – can be used to determine weights of reactants used and products that can be made – Example: CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) 108 Symbols in Chemical Equations  State Symbols (written after the substance’s formula): (g) = gas (↑) 9 (l) = liquid (s) = solid (aq) = aqueous = dissolved in water (ppt) = precipitate (↓)  Energy Symbols (Written above the arrow): high-energy visible – Δ = heat light (hv) – hν = light – shock = mechanical – elec = electrical 109  3.9. Balancing Balancing Chemical Chemical Equations Equations To show the reaction obeys the Law of Conservation of Mass, the equation must be balanced: – We adjust the number of molecules so that there are equal numbers of atoms of each element on both sides of the arrow: CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) 1C + 4H + 4O 1C + 4H + 4O 110 Balancing Chemical Equations: Practice  When aluminum metal reacts with oxygen, it produces a white powdery compound, aluminum oxide aluminum(s) + oxygen(g) aluminum oxide(s) ….. 4 Al(s) + ….. 3 O2(g) ….. 2 Al2O3(s)06 4 4 Al(s) + 3 O2(g) 2 Al2O3(s)  The coefficients required to balance this equation are: 4, 3, 2, respectively. 111 Assessment I 1- Give the coefficients that are necessary to balance each of the following equations: He Cla 2H Cl CUZO C 2 CUT Co z 3 y 2 3 4 3 0 2- What is the coefficient of H2O when each of the following equations are balanced? 2 8 10 112 Types of Chemical Equations  Two types of equations are used to represent chemical reactions: 1- Molecular equations and 2- Ionic equations Example of the molecular equation: NaOH + HCI NaCl + H2O  This equation shows that 1 mole of sodium hydroxide will neutralize exactly 1 mole of hydrochloric acid to form exactly 1 mole of sodium chloride and 1 mole of water; in other words, it represents the exact stoichiometry of the reaction.  A molecular equation is an equation in which the formulas of the compounds are written as though all substances exist as molecules. 113 Types of Chemical Equations  However, there is a better way to show what is happening in this reaction  Sodium hydroxide, hydrochloric acid, and sodium chloride are all strong electrolytes and in solution are 100 % ionized.  Therefore, a solution of sodium hydroxide contains only Na+ and OH- ions and no NaOH molecules. NaOH Na+ + OH-  Similarly, a solution of hydrochloric acid contains only H+ and Cl- ions and no HCI molecules HCI H+ + Cl-  Also, a solution of sodium chloride contains Na+ and Cl- ions and no NaCI molecules. NaCI Na+ + Cl- 114 Types of Chemical Equations  Water, on the other hand, is an extremely weak electrolyte and exists almost entirely in the form of H2O molecules.  So, when a solution of sodium hydroxide is mixed with a solution of hydrochloric acid the expression showing the ionic and molecular species involved would be: Na+ + OH- + H+ + Cl- Na+ + Cl- + H2O  It is clear that both Na+ and Cl- do not altered in the reaction. Therefore, the actual change which takes place is expressed by Net Ionic Equation : H+ + OH- H2O 115 How to write Net Ionic Equation? 1) Write the balanced molecular equation. 2) Write the ionic equation showing the strong electrolytes. 3) Determine if there is any precipitate from the solubility rules. 4) Cancel the similar ions on both sides of the ionic equation. 116 How to write Net Ionic Equation? Example 1: Write the net ionic equation of the reaction of sodium hydroxide with hydrochloric acid.  The molecular equation: NaOH + HCI NaCl + H2O  The ionic equation showing the strong electrolytes: 1 Na+ + OH- + H+ + Cl- Na+ + Cl- + H2O 11  No precipitate is formed; so cancel similar ions: Na+ + OH- + H+ + Cl- Na+ + Cl- + H2O  Net ionic equation will be: H+ + OH- H2O 117 How to write Net Ionic Equation? Example 2: Write the net ionic equation of the reaction of a solution of silver nitrate with hydrochloric acid.  The molecular equation: AgNO3 + HCI AgCl  + HNO3  The ionic equation showing the strong electrolytes: Ag+ + NO3- + H+ + Cl- AgCl + NO3- + H+  Silver chloride (white precipitate) is formed; then cancel the similar ions: Ag+ + NO3- + H+ + Cl- AgCl  + NO3- + H+ I 1  Net ionic equation will be: Ag+ + Cl- AgCl  118 Chemical Bonds Compounds are made of atoms held together by bonds. Chemical bonds are forces of attraction between atoms. The bonding attraction comes from attractions between protons and electrons of bonded atoms. Bonds can form between atoms of the same element, or between atoms of different elements. Chemical bonds form because they lower the potential energy between the charged particles that compose atoms. 119 Types of Chemical Bonds  Chemical bonds can be classified into three types, depending on the types of atoms involved in the bonding:  Ionic bond  Covalent bond Intramolecular Force  Metallic bond 120 Types of Chemical Bonds: The Ionic Bond Ionic bond: results when electrons have been transferred between atoms, resulting in oppositely charged ions that attract each other.  Generally formed when metal atoms bond to nonmetal atoms.  Method: electron transfer. 121 Types of Chemical Bonds: The Covalent Bond Covalent bond: results when two atoms share some of their electrons:  Generally formed when nonmetal atoms bond together  Shared electrons hold the atoms together by attracting nuclei of both atoms.  Method: electron sharing  Multiple Covalent Bonds:  Single covalent bond: A covalent bond formed by sharing one electron pair (2eˉ). Represented by a single line: H−H  Double covalent bond: formed by sharing two electron pairs (4eˉ). Represented by a double line: O=O  Triple covalent bond: formed by sharing three electron pairs (6eˉ). Represented by a triple line: N≡N 122 Types of Chemical Bonds: The Coordinate Bond Coordinate bond (also called a dative covalent bond): is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.  Coordinate bond is a sharing of lone pair of electrons from one atom called donor (Lewis base) to another atom called acceptor (Lewis acid). Mr  Lewis acid: electron pair acceptor e.g. H+, AlCl3, FeBr3, BF3.  Lewis base: electron pair donor e.g. compounds containing heteroatoms (O, S, N) e.g. NH3, H2O. 123 Representing Valence Electrons with Dots (Lewis Structures) Lewis Structures: simple diagrams to visualize the number of valence electrons in atoms of main-group elements by dots. The dots are placed around the element’s symbol with a maximum of two dots per side. Each dot represents one valence electron.  Remember: the number of valence electrons for main group element is equal to the group number of the element (except for helium, which is in group 8A but has only two valence electrons).  Note: While the exact location of dots is not critical, here we first place dots singly before pairing (except for helium which always has two paired dots) 124 Representing Valence Electrons with Dots (Lewis Structures)  The electron configuration of Oxygen is as follows:  Its Lewis structure is as follows: 125 Representing Valence Electrons with Dots (Lewis Structures)  Lewis structure for all period 2 elements: c n Practice: Draw the Lewis dot structure of a phosphorus atom. Solution: Since phosphorus is in Group 5A in the periodic table, it has 5 valence electrons. Represent these as five dots surrounding the symbol for phosphorus: 126 Lewis Structures: For Covalent Bonding Hydrogen and oxygen have the following Lewis structures: In water, hydrogen and oxygen share their valence electrons so that each hydrogen atom gets a duet and the oxygen atom gets an octet. 127 Lewis Theory Predicts That Hydrogen Should Exist as H2 The individual Hydrogen atoms has the following Lewis structure: When two hydrogen atoms share their valence electrons, they each get a duet, a stable configuration for hydrogen. Lewis theory predicts that elemental hydrogen exists as a diatomic molecule (H2). 128 Lewis Structures: Double and Triple Covalent Bonds Oxygen exists as a diatomic molecule (O2): O Nitrogen exists as a diatomic molecule (N2): O s N N Ma 129 Assessment 1. Write the Lewis structure for each atom or ion: a. Al b. sodium ion c. magnesium ion d. chloride ion 2. Use Lewis structures to explain why each element occurs as diatomic molecules: a. hydrogen b. bromine c. oxygen d. nitrogen 3. Write the Lewis structure for each compound: a. PH3 b. SCl2 c. HI d. CH4 e. NaF f. CaO g. SrBr2 h. K2O 4. Determine whether a bond between each pair of atoms would be nonpolar covalent, polar covalent, or ionic. a. Br & Br b. C & Cl c. Mg & I d. Sr & O 5. Order these compounds in order of increasing carbon–carbon bond strength and in order of decreasing carbon–carbon bond length: HC≡CH , H2C═CH2 , H3C─CH3 130  B- Molecules, Compounds and Chemical Bonds 4- Intermolecular Forces and Bond Polarity  Intermolecular Forces 4 55 Intermolecular forces: are the attractive forces that exist between molecules. intra jib inter if it  In contrast to intermolecular forces, intramolecular forces hold atoms together in a molecule.  Generally, intermolecular forces are much weaker than intramolecular forces. It usually requires much less energy to evaporate a liquid than to break the bonds in the molecules of the liquid. 81 I From the chemistry point of view: 64 16 1  The strength of the intermolecular 5 forces (IMF) determines whether D a compound has a high or low melting point and boiling point, and thus if the compound is a solid, liquid, or gas at a given temperature. 2  IMF influences the solubility of substances in various solvents.  IMF affects the rate and outcome of chemical reactions by 3 influencing how reactant molecules come together and interact. 132 65 43 Intermolecular Forces in Covalent Molecules  There are three different types of intermolecular forces in covalent molecules, presented in order of increasing strength: Iiis  London dispersion forces (also called van der Waals forces)  Dipole–dipole interactions (also called van der Waals forces)  Hydrogen bonding i i_ Is was 451 NOF Intramolecular 14 force Intermolecular Forces: London Dispersion Forces London dispersion forces: are very weak interactions due to the momentary changes in electron density in a molecule. Etheric EH Temporary dipole 545 j.ws I All covalent compounds exhibit London dispersion forces. These intermolecular forces are the only intermolecular forces present in nonpolar compounds. The strength of these forces is related to the size of the molecule. v 6 18 of 05 g in a'd  The larger the molecule, the larger the attractive force between two 0 molecules, and the stronger the intermolecular forces. 134 Intermolecular Forces: Dipole–dipole interactions Dipole–dipole interactions: are the attractive forces between the permanent dipoles of two polar molecules. GII  For example, the carbon–oxygen bond in formaldehyde, H2C=O, is polar because oxygen is more electronegative than carbon. This polar bond gives formaldehyde a permanent imbiti dipole, making it a polar molecule.  The dipoles in adjacent formaldehyde molecules can align so that the partial positive and partial negative charges are close to each other. These attractive forces due to permanent dipoles are much stronger than London dispersion forces. Ñm 06.4 aim MW 135  Thestrongest one Intermolecular Forces: Hydrogen Bonding Nof Hydrogen bonding is a strong type of dipole-dipole interaction which occurs when a hydrogen atom bonded to O, N, or F, is electrostatically attracted to an O, N, or F atom in another molecule. stick id Ball  Hydrogen bonding is only possible between two molecules that contain a hydrogen atom bonded to a very electronegative atom—that is, oxygen, nitrogen, or fluorine  These forces are weaker than intramolecular bonds, but are much stronger than other intermolecular forces, causing these compounds to have high boiling points.  Hydrogen bonds are the strongest of the three types of intermolecular forces 136 Intermolecular Forces: Types of Hydrogen Bonding Intermolecular hydrogen bond: Refers to reaction between two same or different molecules. im I  Occurs when hydrogen locates between two electronegative groups (N, S, O). He Hro 00 Intramolecular hydrogen bond: Refers to hydrogen bonds within the same molecule. IS  Stronger than intermolecular hydrogen bonds.  Higher boiling and melting points. Intramolecular hydrogen bonding of salicylaldehyde 137 Intermolecular Forces: Types of Hydrogen Bonding ii i 3 Examples of Intra- and Intermolecular Hydrogen Bonds intra inter Intermolecular Forces: Ion-dipole Forces Ion–dipole forces: are the electrostatic attractions between a charged ion and a dipole.  They are common in solutions and play an important role in the dissolution of ionic compounds, like NaCl, in polar solvent such as water.  The strength of ion–dipole interactions is directly proportional to: 1. the charge on the ion 2. the magnitude of the dipole of polar molecules. 139 Why are intermolecular forces important in pharmacy? 20 IMF forces are of great importance in the field of pharmacy, as they play a crucial role in various aspects of pharmaceutical science and drug development. Here are some key ways in which IMF are important in pharmacy: 1- Drug solubility and formulation:  The solubility of a drug molecule in certain solvent is largely determined by the IMF between the drug and the solvent. Understanding these forces is crucial for designing effective drug formulations. 2- Drug absorption and bioavailability:  IMF influence the ability of a drug to be absorbed from the gastrointestinal tract into the bloodstream. For example, hydrogen bonding and ionic interactions can affect the permeability of drug molecules across biological membranes.  The bioavailability of a drug, which is the fraction of the administered dose that reaches the systemic circulation, is also affected by intermolecular forces during absorption, distribution, and metabolism. Why are intermolecular forces important in pharmacy? 3- Drug stability and shelf life: If I IMF such as hydrogen bonding and van der Waals interactions, can stabilize the structure of drug molecules and influence their resistance to degradation. Understanding the IMF involved in drug-excipient and drug-drug interactions is crucial for developing stable pharmaceutical formulations with an adequate shelf life. 4- Drug delivery and targeting: IMF play a role in the design of drug delivery systems, such as liposomes, nanoparticles, and polymeric carriers, which aim to improve the targeting and controlled release of drugs. 5- Protein-drug interactions: Many drugs exert their therapeutic effects by interacting with specific proteins, such as enzymes, receptors, or transport proteins. The strength and nature of these protein-drug interactions are determined by IMF. Intermolecular Forces: Problems 1- What types of intermolecular forces are present in each compound: (a) HCl; (b) C2H6 (ethane); (c) NH3? N.B. London dispersion forces are present in all covalent compounds. Dipole–dipole interactions are present only in polar compounds with a permanent dipole. Hydrogen bonding occurs only in compounds that contain an O – H, N – H, or H – F bond. a. HCl has London forces like all covalent compounds. HCl has a polar bond, so it exhibits dipole–dipole interactions. HCl has no H atom on an O, N, or F, so it has no intermolecular hydrogen bonding. b. C2H6 is a nonpolar molecule since it has only nonpolar C – C and C – H bonds. Thus, it exhibits only London forces. c. NH3 has London forces like all covalent compounds. NH3 has a net dipole from its three polar bonds, so it exhibits dipole–dipole interactions. NH3 has a H atom bonded to N, so it exhibits intermolecular hydrogen bonding.142 Intermolecular Forces: Problems Assessment 2. What types of intermolecular forces are present in each molecule? a. Cl2 b. HCN c. HF d. CH3Cl e. H2 3. Which of the compounds in each pair has stronger intermolecular forces? a. CO2 or H2O b. CO2 or HBr c. HBr or H2O d. CH4 or C2H6B H2o HBV H2o 246 hydrogen bonding: 4. Determine which compound can form inter or intramolecular a. 7.30 intra inter intra inter e e 143 Electronegativity and Bond Polarity  Electronegativity (EN): is the ability of an atom (in a molecule) to 1 attract the bond electrons to itself. IN MG NTS  is higher for nonmetals; and lower for metals  The greater the difference in electronegativity (ΔEN), the more polar the bond. δ+ and δ- in polar molecular compounds represent the partial positive and negative charges, to differentiate them from the full charge I (+ or -) on ions in ionic compounds. w.es who I if 144 I Electronegativity Values for Elements (Unitless) 43hr5538 5 2T 145  Electronegativity and Bond Types Covalent Polar I 12 Hc Ionic Nad range Eiii.tn em 146 Electronegativity and Bond Types: Examples Assessment - Example: Based on the of values of electronegativity (EN) of elements, which bond is more polar: (B-Cl) or (C-Cl)? Answer: - The ΔEN of Cl and B = 3.0 - 2.0 = 1.0 CovalentPolar _f - The ΔEN of Cl and C = 3.0 - 2.5 = 0.5 covalent Polan  Hence, the B-Cl bond is more polar. - Exercise 1: Which of the following bonds is the most polar? (a) HF (b) SeF (c) NP (d) GaCl - Exercise 2: Predict the type of each bond (use the table of EN values): (a) HBr (b) OO (c) HO (d) SO 147