Molecular Geometry PDF

MOLECULAR GEOMETRY ANNE JLAINE A. LLOSA, SST-I  OBJECTIVES:  Explain the properties of covalent molecular compounds in terms of their structure  Describe the geometry of simple compounds  Determine the polarity of simple molecules MOLECULAR GEOMETRY  The three-dime...

MOLECULAR GEOMETRY ANNE JLAINE A. LLOSA, SST-I  OBJECTIVES:  Explain the properties of covalent molecular compounds in terms of their structure  Describe the geometry of simple compounds  Determine the polarity of simple molecules MOLECULAR GEOMETRY  The three-dimensional arrangement oof atoms in a molecule.  A molecule’s geometry affects its physical and chemical properties, such as melting point, boiling point, density, and the types of reactions it undergoes. MOLECULAR GEOMETRY  Canbe determined by experiment such as x-ray diffraction  However, geometry of simple molecules can be predicted without experimentation by knowing the number of electron surrounding a central atom in its Lewis structure.  BASIS  The prediction rests on the assumption that all electron pairs in the valence shell around a central atom repel one another and want to be as far apart from one another as possible.  They assume a geometry or orientation that will minimize the repulsions; this is the stable orientation and the one with lowest energy VSEPR THEORY VALENCE-SHELL ELECTRON-PAIR REPULSION THEORY This Photo by Unknown Author is licensed under CC BY-SA KEY IDEAS OF THE THEORY  Electron pairs stay as far apart from each other as possible to minimize repulsions  Molecular shape is determined by the number of bond pairs and lone pairs around the central atom  Multiple bonds are treated as if they were in single bonds (in making the prediction)  Lone pairs occupy more volume than bond pairs. Lone pair – lone pair repulsions are greater than lone pair – bond pair repulsions which in turn are greater than bond pair – bond pair repulsions POLARITY OF MOLECULES Factors that determine the polarity of molecules:  Thepolarity of bonds between atoms which can be studied based on electronegativity  Thegeometrical shape of the molecule which can be predicted via the VSEPR theory Non-Polar Covalent Bonds  Occur when electron pairs are shared equally  The difference in electronegativity between atoms is less than 0.5 Polar Covalent Bonds  Occur when electron pairs are unequally shared  The difference in electronegativity between atoms is significant (0.5- 2.0)  DIPOLE MOMENT  The polarity of a bond can be experimentally measured in terms of the dipole moment, μ  It is the product of the charge, Q, and the distance between the charges, r.  To maintain neutrality, the changes on the ends of the molecule must be equal in magnitude but opposite in sign. Nonpolar molecules have no net dipole moment. Polar molecules exhibit dipole moments.  In the presence of an electric field, the positive end of the molecules orient themselves towards the negative plate. POLAR OR NONPOLAR? BF3 POLAR OR NONPOLAR? CH2Cl2 Activity: 1 whole yellow pad Draw the Lewis structure of the following molecules and determine its molecular geometry and polarity. 1. PI3 6. HCl 2. N2 7. FCN 3. H 2O 8. HNO 4. AsBr3 9. AsN 5. SiCl4 10. CH3NH2 GENERAL CHEMISTRY II Anne Jlaine A. Llosa, SST-I REVIEW Br2 NH3 C H 3B r Propanone KINETIC MOLECULAR THEORY All matter is made of tiny particles These particles are in constant motion The speed of particles is proportional to temperature. Increased temperature means greater speed. Solids, liquids, and gases differ in distances between particles, in the freedom of motion of particles and in the extent to which the particles interact. Properties of Matter Molecular Behavior Gas Liquid Solid Volume Shape Density Compressibility Motion of Molecules KINETIC MOLECULAR VIEW OF THE THREE S TAT E S ATTRACTIVE FORCES VS. KINETIC ENERGY GASES Attractive forces are weak relative to kinetic energy Attractive forces are stronger because particles LIQUID have less kinetic energy Attractions dominate motion. SOLID Particles are fixed in place relative to each other W H AT I S A P H A S E ? W H AT I S A P H A S E ? is a homogeneous part of the system in contact with other parts of the system but separated from them by a well-defined boundary. PHASE CHANGES  WHAT HOLDS THE PARTICLES IN THE SOLID AND LIQUID STATES? INTERMOLECULAR FORCES or INTRAMOLECULAR FORCES INTERMOLECULAR FORCES INTRAMOLECULAR FORCES Are attractive forces between Are attractive forces within a molecules molecule 41 kJ is needed to vaporize 1 930 kJ is required to break all O- mole of water H bonds in 1 mole of water Generally, intermolecular forces are much weaker than intramolecular forces. INTRAMOLECULAR FORCES OF AT TRACTI O N INTERMOLECULAR FORCES OF AT TRACTION 1. LONDON DISPERSION FORCES (LDF) Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules ion-induced dipole interaction dipole-induced dipole interaction  These forces are present between all types of molecules due to the movement of electrons. As electrons move around the nucleus, an uneven distribution causes momentary charge separations. Slightly positive sides of a molecule are attracted to the slightly negative sides of the adjacent molecule.  The dipole in the atom (or nonpolar molecule) is said to be an induced dipole because the separation of positive and negative charges in the atom (or nonpolar molecule) is due to the proximity of an ion or a polar molecule. POLARIZABILITY - is the ease with which the electron distribution in the atom or molecule can be distorted Polarizability increases with: ✓Greater number of electrons ✓More diffused electron cloud iso-Pentane, bp = 27.8°C EXAMPLE: F2, the lightest halogen, is a gas, Br2 is a liquid, and the heavier I2,is a solid at room conditions. Further, the more atoms that make up the molecules, the stronger are the dispersion forces. 2. DIPOLE-DIPOLE FORCES Attractive forces between polar molecules Orientation of Polar Molecules in a Solid 3. ION-DIPOLE FORCES Attractive forces between an ion and a polar molecule Ion-Dipole Interaction Interaction Between Water and Cations in solution 33 4. HYDROGEN BOND is a special dipole-dipole interaction between they hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A H…B or A H…A A & B are N, O, or F HYDROGEN BONDING IN DNA “Measure” of intermolecular force boiling point melting point Hvap Hfus Hsub INTERMOLECULAR FORCES OF AT TRACTI O N SUMMARY COMPREHENSION CHECK: Determine what type/s of intermolecular forces exist between each of the following molecules? HBr CH4 SO2 C H 3C H 2O H CO2 H 2O a n d H F H 2 and CO 2 H 2 and NaF H 2O a n d K C l ARRANGE ACCORDING TO INCREASING IMF STRENGTH N2 HCN HF ARRANGE ACCORDING TO INCREASING IMF STRENGTH Water Methanol Ethanol Diethyl ether ARRANGE IN INCREASING BOILING POINT HCl HBr HF HI INTERMOLECULAR FORCES OF ATTRACTION ANNE JLAINE A. LLOSA, RCh SST-I REVIEW: DETERMINE THE TYPE/S OF INTERMOLECULAR FORCE/S EXISTING BETWEEN THE FOLLOWING PAIRS (a) HBr and H2S (b) Cl2 and CBr4 (c) I2 and NO3 2- (d) NH3 and C6H6 (e) H2O and CH3CH2OH WHY IS THE HYDROGEN BOND CONSIDERED A “SPECIAL” DIPOLE-DIPOLE INTERACTION?  PROPERTIES OF LIQUIDS: Describe the properties of liquids: surface tension, viscosity, vapor pressure, boiling point, and molar heat of vaporization; Explain the effect of intermolecular forces on these properties; and Relate the properties of water to intermolecular forces that operate among its molecules DEFINITION OF TERMS: FLUID A gas or a liquid; a substance that can flow DEFINITION OF TERMS: SURFACE TENSION The measure of the elastic force in the surface of a liquid. It is the amount of energy required to stretch or increase the surface of a liquid by a unit area DEFINITION OF TERMS: C APILLARY ACTION The tendency of a liquid to rise in narrow tubes or to be drawn into small openings DEFINITION OF TERMS: VISCOSITY A measure of a fluid’s resistance to flow DEFINITION OF TERMS: VAPOR A gaseous substance that exist naturally as a liquid or solid at normal temperature DEFINITION OF TERMS: VAPORIZATION The change of phase from liquid to gas DEFINITION OF TERMS: VAPOR PRESSURE OF A LIQUID The equilibrium pressure of a vapor above its liquid; that is, the pressure exerted by the vapor above the surface of the liquid in a closed container DEFINITION OF TERMS: BOILING POINT The temperature at which a liquid boils. The boiling of a liquid when the external pressure is 1 atm is called the normal boiling point.  DEFINITION OF TERMS: MOLAR HEAT OF VAPORIZATION The energy required to vaporize 1 mole of a liquid at a given temperature PROPERTIES OF LIQUIDS  SURFACE TENSION - measure of the elastic force in the surface of a liquid - the amount of energy required to stretch or increase the surface of a liquid by a unit area  Strong intermolecular forces High surface tension SURFACE TENSION  WATER STRIDER WALKING ON THE SURFACE OF A QUIET POND  CAPILLARY ACTION -is the tendency of a liquid to rise in narrow tubes or be drawn into small openings such as those between grains of a rock - is a result of intermolecular forces of attraction between the liquid and solid materials ADHESION is an COHESION is the attraction between intermolecular attraction unlike molecules between like molecules T WO T Y P E S O F F O R C E S I N VO LV E D I N C A P I L L A RY AC T I O N  When the When cohesive cohesive forces forces between between the the liquid liquid molecules molecules are are lesser than greater than the the adhesive adhesive forces forces between between the the liquid and liquid and the the walls of the walls of the container, the container, the surface of the surface of the liquid is liquid is convex concave  VISCOSITY Is a measure of a fluid’s resistance to flow  GLYCEROL G IV E N T HE MOLE C U LAR ST RU CTURES O F WAT E R A N D G LY C E RO L , C A N YO U T E L L W H Y G LY C E RO L H A S A H I G H E R V I S C O S I T Y T H A N WAT E R ? WATER ALL THE SUBSTANCES IN THE LIST ARE HYDROC ARBONS AND NONPOLAR, WHAT C AUSES THE VISCOSITIES OF THE HYDROC ARBONS ON THE LIST? VAPOR PRESSURE OF A LIQUID  VAPOR PRESSURE OF THE LIQUID. The pressure exerted by the gas in equilibrium with a liquid in a closed container at a given temperature The equilibrium vapor pressure is the maximum vapor pressure of a liquid at a given temperature and that it is constant at a constant temperature. It increases with temperature.  EQUILIBRIUM VAPOR PRESSURE - is the vapor pressure measured when a dynamic Dynamic Equilibrium equilibrium exists Rate of Rate of = evaporation condensation between condensation and evaporation H2O (l) H2O (g)  When the rate of condensation of the gas becomes equal to the rate of evaporation of the liquid, the gas in the container is said to be in equilibrium with the liquid. liquid ⇋ vapor (gas) In this condition, the amount of gas and liquid no longer changes. H2O (l) H2O (g)  Measurement of Vapor Pressure Before At Evaporation Equilibrium 30  VAPOR OF FOUR COMMON LIQUIDS, SHOWN AS A FUNCTION OF TEMPERATURE VAPOR PRESSURE VS. INTERMOLECULAR FORCE When liquids evaporate, the molecules have to have sufficient energy to break the attractive forces that hold them in the liquid state. The stronger these intermolecular forces are, the greater the amount of energy needed to break them  Molar heat of vaporization (Hvap) is the energy required to vaporize 1 mole of a liquid at its boiling point. Clausius-Clapeyron Equation Hvap P = (equilibrium) vapor pressure ln P = - +C T = temperature (K) RT R = gas constant (8.314 J/K mol) Vapor Pressure Versus Temperature THE RELATIONSHIP BETWEEN VAPOR PRESSURE AND STRENGTH OF INTERMOLECULAR FORCES IS CONSISTENT WITH THE TRENDS IN TWO OTHER PROPERTIES OF LIQUIDS, THE ENTHALPY OR MOLAR HEAT OF VAPORIZATION, AND THE BOILING POINT OF THE LIQUID.  The boiling point is the temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure. The normal boiling point is the temperature at 3 which a liquid boils when 7 the external pressure is 1 atm.  PRACTICE At 50° C the vapor pressure of ethanol is 0.30 atm, acetic acid is 0.08 atm, water is 0.12 atm, and acetone is 0.84 atm. A. Arrange these substances in order of increasing rates of evaporation. B. Arrange these substances in order of increasing boiling point temperature. C. Arrange these substances in order of increasing intermolecular forces. 38  ACTIVITY (1/2 CROSSWISE) A. Identify if the molecules are polar, non-polar, or ionic and determine the strongest intermolecular force present in between the following molecules: 1. H2S 4. H2O and BF2Cl 2. CH3COOH 5. PCl5 and BF3 3. CS2 6. CH4 and Na+ B. Arrange the following molecules according to increasing vapor pressure: Kr I2 CO2 H 2O CH2Cl2 39  SOLIDS Anne Jlaine A. Llosa, RCh SST-I How do solids behave? Do you think that the structure of solids affect its properties? If yes, how? Why is diamond stronger than graphite? 4 OBJECTIVES: Describe the difference in structure of crystalline and amorphous solids Describe the different types of crystals and their properties: ionic, covalent, molecular, and metallic 5 TWO GENERAL TYPES OF SOLIDS CRYSTALLINE SOLID AMORPHOUS SOLID - possesses rigid and - does not possess a well- long-range order. In a defined arrangement and crystalline solid, atoms, long range molecular molecules, or ions order occupy specific (predictable) positions 6 SMALL GROUP DISCUSSION Divide into 4 groups Analyze the pictures given Answer the given questions Pick two representatives to present your answers 7 DIRECTIONS Enumerate at least 2 differences among the given pictures Identify the 2 types of solids given to you 8 FOCUS QUESTIONS: A. What are the two general types of solids? What features can be used to distinguish a crystalline solid from an amorphous solid? B. What is the distinguishing feature of crystalline solids? How are the structures of crystals determined? C. What are the four types of crystals? What forces bind the unit particles of each type of crystal? What are the properties of each type of crystal? 9 A. WHAT ARE THE TWO GENERAL TYPES OF SOLIDS? WHAT FEATURES CAN BE USED TO DISTINGUISH A CRYSTALLINE SOLID FROM AN AMORPHOUS SOLID? Solids can be categorized into two groups: the crystalline solids and the amorphous solids Their differences in properties arise from the presence or absence of long range order of arrangements of the particles in the solid 10  1. ARRANGEMENT OF PARTICLES The compounds 1. They can form a regular repeating of a solid can be three-dimensional structure called a arranged in two crystal lattice, thus producing a crystalline solid, or general ways: 2. They can aggregate with no particular long range order, and form an amorphous solid 11 CRYSTALLINE SOLIDS - are arranged in fixed geometric patterns or lattices The ordered arrangement of their units maximizes the space they occupy and are essentially incompressible 12 SEVEN BASIC UNIT CELLS 13 THREE TYPES OF CUBIC UNIT CELLS 14 EXAMPLES OF CRYSTALLINE SOLIDS 15 AMORPHOUS SOLIDS - have a random orientation of particles They are considered super-cooled liquids where molecules are arranged in a random manner similar to the liquid state 16 EXAMPLES OF AMORPHOUS SOLIDS 17  More than 90% of naturally occurring and artificially prepared solids are crystalline. Minerals, sand, clay, limestone, metals, alloys, carbon (diamond and graphite), salts (e.g. NaCl and MgSO4), all have crystalline structures Amorphous solids (e.g. glass), like liquids, do not have long range order, but may have a limited, localized order in their structures. 18 2. BEHAVIOR WHEN HEATED The presence or absence of long- range order in the structure of solids results in a difference in the behavior of the solid when heated 19  ❑ The structures of crystalline solids are built from repeating units called crystal lattices. The surroundings of particles in the structure are uniform, and the attractive forces experienced by the particles are of similar type and strength 20  ❑ These attractive forces are broken by the same amount of energy, and thus, crystals become liquids at a specific temperature ❑ At this temperature, physical properties of crystalline solids change sharply 21 ❑Amorphous solids soften gradually when they are heated ❑They tend to melt over a wide range of temperature ❑This behavior is a result of the variation in the arrangement of particles in their structures, causing some parts of the solid to melt ahead of other parts B. WHAT IS THE DISTINGUISHING FEATURE OF CRYSTALLINE SOLIDS? HOW ARE THE STRUCTURES OF CRYSTALS DETERMINED? 23 WHAT IS THE DISTINGUISHING FEATURE OF CRYSTALLINE SOLIDS? Crystal lattice crystalline solids are characterized by a regular repeating structure called the crystal lattice. 24 HOW ARE THE STRUCTURES OF CRYSTALS DETERMINED? X-ray Diffraction is a technique used to determine the atomic and molecular structure of a crystal, wherein atoms cause beams of incident rays to diffract into many specific directions 25 HOW ARE THE STRUCTURES OF CRYSTALS DETERMINED? The scattered rays interfere with each other and produce a pattern of spots of different intensities that can be recorded on film, such as that shown in the figure below. 26 C. WHAT ARE THE FOUR TYPES OF CRYSTALS? WHAT FORM OF UNIT ARTICLES MAKES UP EACH TYPE OF CRYSTAL? WHAT FORCES BIND THE UNIT PARTICLES OF EACH TYPE OF CRYSTAL? WHAT ARE THE PROPERTIES OF EACH TYPE OF CRYSTAL? 27 TYPES OF CRYSTALS They differ in the kind of articles that make up the crystal and the attractive forces that hold these particles together 1. Metallic crystals 2. Ionic crystals 3. Molecular crystals 4. Covalent Network crystals 28 1. METALLIC CRYSTALS are made of atoms that readily lose electrons to form positive ions (cations), but no atoms in the crystal would readily gain electrons. The metal atoms give up their electrons to the whole crystal, creating a structure made up of an orderly arrangement of cations surrounded by delocalized electrons that move around the crystal. 29  The crystal is held together by electrostatic interactions between the cations and delocalized electron. These interactions are called metallic bonds. This model of metallic bonding is called the “sea of electrons” model. 30 31 32 33 PROPERTIES OF METALLIC CRYSTALS 34 2. IONIC CRYSTALS Ionic crystals are made of ions (cations and anions). These ions form strong electrostatic interactions that hold the crystal lattice together. The electrostatic attractions are numerous and extend throughout the crystal since each ion is surrounded by several ions of opposite charge, making ionic crystals hard and of high melting points. 35 Ionic Crystals Lattice points occupied by cations and anions Held together by electrostatic attraction Hard, brittle, high melting point Poor conductor of heat and electricity CsCl ZnS CaF2 36 37 38 38 THE ENERGY NEEDED TO BREAK THE CRYSTAL OF IONIC SUBSTANCES WILL DEPEND ON: magnitude of charges on the ions - (the 2+ and 2- ions attract each other stronger in MgO than 1+ and 1- in NaCl) sizes of the ions - (attractions are less between the bigger ions in RbI and as such less heat energy is needed to separate them than the smaller ions in NaCl). 39 PROPERTIES OF IONIC CRYSTALS 40  Ionic substances can conduct electricity in the liquid or molten state or when dissolved in water, indicating that in these states, charged particles are able to move and carry electricity. However, the solid state is generally nonconducting since the ions are in fixed positions in the crystal lattice and are unable to move from one point to another 41  Ionic crystals are brittle, and would shatter into small pieces when deformed or when pressure is applied on the crystal. The shifting of ions cause repulsions between particles of like charges. 42 3. MOLECULAR CRYSTALS ▪ The atoms or molecules are held together by a mix of hydrogen bonding/dipole-dipole and dispersion forces, and these are the attractive forces that are broken when the crystal melts.. ▪ most molecular crystals have relatively low melting points. ▪ such as in noble gases, or molecules, such as in sugar, C12H22O11, iodine, I2, and naphthalene, C10H8. 43 43 44 Molecular Crystals Lattice points occupied by molecules Held together by intermolecular forces Soft, low melting point Poor conductor of heat and electricity water benzene 45 PROPERTIES OF MOLECULAR CRYSTALS 46 47  4. COVALENT NETWORK CRYSTALS Covalent network crystals are made of atoms in which each atom is covalently bonded to its nearest neighbors. The atoms can be made of one type of atom (e.g. C- diamond and C-graphite) or can be made of different atoms (e.g. SiO2 and BN). In a network solid, there are no individual molecules and the entire crystal may be considered one very large molecule. 48 49 50  The valence electrons of the atoms in the crystal are all used to form covalent bonds. Because there are no delocalized electrons, covalent network solids do not conduct electricity. Covalent bonds are the only type of attractive force between atoms in the network solid. Rearranging or breaking of covalent bonds requires large amounts of energy; therefore, covalent network solids have high melting points. 51  Covalent bonds are extremely strong, so covalent network solids are very hard. Generally, these solids are insoluble in water due to the difficulty of solvating very large molecules. Diamond is the hardest material known, while cubic boron nitride (BN) is the second-hardest. Silicon carbide (SiC) is very structurally complex and has at least 70 crystalline forms. 52 PROPERTIES OF COVALENT NETWORK SOLIDS 53 54 55  An amorphous solid does not possess a well-defined arrangement and long-range molecular order. A glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing Crystalline Non-crystalline 56 quartz (SiO2) quartz glass PHASE CHANGES Anne Jlaine A. Llosa, RCh Special Science Teacher I What phase/s of matter exist in the following image? What phase/s of matter exist in the following image? What phase/s of matter exist in the following image? What are other examples of phase changes?  OBJECTIVES: ❑ Describe the nature of the following changes in terms of PHASE energy change and the increase or CHANGES decrease in molecular order: solid- liquid, liquid-vapor, and solid-vapor ❑Interpret the phase diagram of water and carbon dioxide What causes the phase change in matter?  are transformations of matter from one physical state to another they occur when energy (usually in the form of heat) is added or removed from Phase Change a substance characterized by changes in molecular order - molecules in the solid phase have the greatest order, while those in the gas have the greatest randomness or disorder Difference in molecular order of a substance in the solid, liquid, and gaseous states Types of Endothermic processes Phase Exothermic processes Changes Endothermic processes ✓Melting - these changes take place ✓Vaporization when heat is absorbed (heat ✓Sublimation gained) Exothermic processes ✓Freezing - these changes take place ✓Condensation when heat is released (heat ✓Deposition lost) Figure 3: The different changes in state that matter undergoes (Image Source: http:// www.shmoop.com/matter-properties/test-your-knowledge.html) Phase Changes Least Order Greatest 14 Order How does a change in energy affect phase changes?  1. The added heat increases the kinetic When a energy of the particles and the particles substance is move faster accompanied by an increase heated, the in temperature. added energy 2. The added heat is used to break is used by the attractive forces between particles. There substance in is no observes increase in temperature either of two when this happens. Often a change in ways: physical appearance of the substance is observes, such as a phase change.  The change in temperature of a substance as it is being heated can be shown in a graph called the heating curve.  1. A decrease in the kinetic energy of When a the particles. The motion of substance particles slow down, and a undergoes a decrease in temperature is removal or observed. release of 2. Forces of attraction are formed, heat, it results and a phase change may occur. No in two ways: change in temperature is observed. Calculations A. Heat The amount of heat received or change with removed from the sample given a change change in in temperature can be calculated using temperature specific heat of substance  - amount of heat needed to raise the temperature of 1 gram of a substance by 1⁰C. Specific Heat - also equal to the amount of heat lost by 1 gram of substance when its temperature drops by 1 ⁰C.  The specific heat of a substance differs for the solid, liquid, and gaseous states Water as an example, has the following specific heat at different phases: H2O(l) = 4.18 J / g ⁰C H2O(s) = 2.06 J / g ⁰C Heat Change H2O(g) = 2.02 J / g ⁰C with Change in The heat change (q) for this process is given by: Temperature q = m S ΔT where: m = mass of sample in grams S = specific heat of the sample in the appropriate physical state T = change in temperature  You found a piece of copper metal weighing 3.10 g imbedded in an ice block. How much heat is absorbed by the piece of metal as it warms in your Sample hand from the temperature of the ice Problem 1 block at 1.5 ⁰C to your body temperature of 37.0 ⁰C? The specific heat of copper is 0.385 J/g ⁰ C. Assume that the metal is pure copper. Problems Involving Changes of State  Problem: How much energy is required to change 2600 gram of ice at 0˚C into water at the same temperature? Sample Solution: Since the problem indicates no change in temperature and involves a solid phase, then Problem 1 the formula to be used is q = m ΔHfus. q = m ΔHfus = (2600 g) (6.01 kJ/mol) = 867.63 kJ  Problem: How much energy is required to change 2600 gram of water at 100˚C into steam at the same temperature? Sample Solution: Since the problem indicates no change in temperature and involves a liquid phase, then the Problem 2 formula to be used is q = m ΔHvap q = m ΔHvap = (2600 g) (40.79 kJ/mol) = 5,888.62 kJ  Calculate the amount of energy (in kJ) needed to heat 346 gram of liquid water from 0⁰C to 182⁰C. Assume that the Sample specific heat of water is 4.184 J/g ⁰C over Problem 3 the entire liquid range and the specific heat of steam is 1.99 J/g ⁰C. Δhvap of water is 40.79 kJ/mol.  The change in temperature of a substance as it is being heated can be shown in a graph called the heating curve. How can this effect be achieved using CO2 or dry ice? Because carbon dioxide cannot exist as a liquid at atmospheric pressure, the dry ice sublimates and instantly produces a gas, condensing water vapor, and creating a thick white fog. PHASE DIAGRAMS  - a graphical representation of the physical states of a substance under different conditions of temperature and pressure. Phase - It gives the possible combinations of Diagram pressure and temperature at which certain physical state or states a substance would be observed. - Each substance has its own phase diagram.  a. The three areas (SOLID, LIQUID, GAS) b. Three lines (curves) Features of a > melting/freezing curve, Phase > vaporization/condensation curve, Diagram > sublimation/deposition curve c. Two important points > triple point > critical point A. Three areas (SOLID, LIQUID, GAS) B. Three Lines/Curves Melting/ Freezing curve Vaporization/ condensation curve Sublimation/Deposition curve Liquid vapor Solid vapor C. Two Important Points a. Triple point - is the combination of pressure and temperature at which all three phases of matter are at equilibrium. - It is the point on a phase diagram at which the three states of matter coexist. C. Two Important Points b. The critical point - The critical point terminates the liquid/gas phase line. - It is the set of temperature and pressure on a phase diagram where the liquid and gaseous phases of a substance merge together into a single phase.  C. Two Important Points b. The critical point - Beyond the temperature of the critical point, the merged single phase is known as a supercritical fluid. - The temperature and pressure corresponding to this are known as the critical temperature and critical pressure.  This single phase is called a supercritical fluid, which exhibits many of the properties of a gas but has a density more typical of a liquid.  Phase Diagram of Water (H2O)  The solid-liquid equilibrium line (the melting point curve) slopes backwards rather than forwards.  the melting point gets lower at higher pressures. because solid ice is less dense than liquid water caused by the crystal structure of the solid phase.  In the solid forms of water and some other substances, the molecules crystallize in a lattice with greater average space between molecules, thus resulting in a solid occupying a larger volume and consequently with a lower density than the liquid. When it melts, the liquid water formed occupies a smaller volume The normal melting point of water is 273 K, and its normal boiling point is 373 K.  The Phase Diagram for Carbon Dioxide - only thing special about this phase diagram is the position of the triple point, which is well above atmospheric pressure. - It is impossible to get any liquid carbon dioxide at pressures less than 5.2 atmospheres. - At 1 atm pressure, carbon dioxide will sublime at a temperature of 197.5 K (-75.5 °C). - solid carbon dioxide is often known as "dry ice.“ - There is no liquid carbon dioxide under normal conditions - only the solid or the vapor. Interpreting the phase diagram 1. In what phase is the substance at 50 °C and 1 atm pressure? 2. At what pressure and temperature conditions will all three phases of the substance be present? 3. What is the normal melting point of the substance? 4. What phase(s) will exist at 1 atm and 70 °C? 1. List the phase changes a sample of ice would go through if heated to its critical temperature at 1 atm? 2. In what phase does water exist at its triple point? 3. How does the melting point of water change as the pressure increases from 1 atm? SOLUTIONS Anne Jlaine A. Llosa, RCh Special Science Teacher I Objectives: ❑ Describe the different types of solutions ❑ Use different ways of expressing concentration of solutions: percent by mass, mole fraction, molarity, molality, percent by volume, ppm  TYPES OF SOLUTIONS Solution - a homogenous mixture which mainly comprises two components namely the solute and solvent SOLVENT – any substance, usually liquid, which is capable of dissolving one or several substances, thus creating a solution SOLUTE – the substance that is being dissolved to create a solution SOLUTE SOLVENT STATE OF EXAMPLE RESULTING SOLUTION Gas Gas Gas Air Gas Liquid Liquid Soda water Gas Solid Solid H2 gas in Pd Liquid Liquid Liquid Ethanol in water Solid Liquid Liquid NaCl in water Solid Solid Solid Brass (Cu/Zn) Types of solutions depending upon the dissolution of the solute 1. Saturated solution - contains the maximum amount of a solute that will dissolve in a given solvent at a specific temperature 2. Unsaturated solution - contains less solute than it has the capacity to dissolve 3. Supersaturated solution - contains more solute than is present in a saturated solution - not very stable; in time, some of the solute will come out of a supersaturated solution as crystals **Crystallization – the process in which dissolved solute comes out of solution and forms crystals Types of solutions: depending on whether the solvent is water or not Aqueous solution – when a solute is dissolved in water Non-aqueous solution – when a solute is dissolved in a solvent other than water Types of solutions: depending on the amount of solute added to the solvent Dilute solution – contains a small amount of solute in a large amount of solvent Concentrated solution – contains a large amount of solute dissolved in a small amount of solvent Concentration Units Concentration – general measurement unit that reports the amount of solute present in a known amount solution Percent by mass - also called percent by weight or weight percent - the ratio of the mass of a solute to the mass of the solution by 100 percent Example 1: A sample of 0.892 g of potassium chloride (KCl) is dissolved in 54.6 g of water. What is the percent by mass of KCl in the solution? Example 2: A sample of 6.44 g of naphthalene (C10H8) is dissolved in 80.1 g of benzene (C6H6). Calculate the percent by mass of naphthalene in this solution. Mole fraction (X) - the number of moles of a specific component in the solution divided by the total number of moles in the given solution Example 1: 0.100 mole of NaCl is dissolved into 100 g of pure water. What is the mole fraction of NaCl? Example 2: A solution is prepared by mixing 25.0 g of water, and 25.0 g of ethanol. Determine the mole fractions of each substance. Molarity (M) - defined as the number of moles of solute per liter of solution Example 1: What is the molarity of a solution that was prepared by dissolving 14.2 g sodium nitrate in enough water to make 350 mL of solution? Example 2: How many grams of sodium bromide would be needed to prepare 700 mL of 0.230 M NaBr solution? Molality (m) - the number of moles of solute dissolved in 1 kg of solvent Example 1: Calculate the molality of a sulfuric acid solution containing 24.4 g of sulfuric acid in 198 g of water. The molar mass of sulfuric acid is 98.09 g Example 2: What is the molality of a solution containing 7.78 g of urea [(NH2)2CO) in 203 g of water? Percent by volume - the volume of solute divided by the total volume of the solution, multiplied by 100 Example 1: Determine the percent by volume concentration of a solution made by combining 25 mL of ethanol with enough water to produce 200mL of solution? Example 2: The deionization reaction of NaOH with water gives sodium ions and hydroxide ions as shown in the reaction: NaOH + H2O → Na+ + OH- Assume that 6 mL of NaOH is dissolved in 80 mL of water to perform deionization of the basic salt. Determine the percent by volume of NaOH in the solution.  - the number of units of mass of a solute Parts per million per million units of total solution (ppm) - convenient for reporting amount of solute that is present in trace quantities Example 1: What is the concentration of a solution, in ppm, if 0.02 g of NaCl is dissolved in 1000 g of solution? Example 2: If 0.025 g of Pb(NO3)2 is dissolved in 100 g of water, what is the concentration of the resulting solution, in ppm? Comparison of Concentration Units Problem 1: Express the concentration of a 0.396 m glucose (C6H12O6) solution in molarity (M) given that the density of the solution is 1.16 g/mL Problem 2. The density of a 2.45 M aqueous solution of methanol (CH3OH) is 0.976 g/mL. What is the molality of the solution? Problem 3. Suppose you are given a concentrated solution of HCl which is known to be 37.0% HCl (w/w) and has a solution density of 1.19 g/mL. What is the molarity, molality and mole fraction of HCl? Problem 4. What is the mass percent of sucrose in a solution obtained by mixing 225 g of an aqueous solution that is 6.25% sucrose by mass with 135 g of an aqueous solution that is 8.20% sucrose by mass? Problem 5. What mass of a 4.00% NaOH solution by mass contains 15.0 g of NaOH? Activity: 1 whole sheet of paper 1. What is the molality of a solution prepared by dissolving 225 mg of glucose in 5.00mL of ethanol? Density of ethanol is 0.789g/mL. 2. A brand of wine has the alcohol (ethanol) content clearly labelled as 13% v/v. Given that the density of ethanol at room temperature is 0.7892 g/mL, calculate the molarity of ethanol. 3. Which solution is more concentrated: a solution made of 4.5 mol NaOH in 150 mL water or a solution made of 10.0 mol NaOH in 500 mL water? 4. A solution is prepared by dissolving naphthalene (128.17 g/mol) in cyclohexane (84.16g/mol), and the resulting concentration is 14.35% naphthalene by mass. Determine the mole fraction and the molality of the solution. 5. What volume of 0.0995 M Al(NO3)3 will react with 3.66 g of Ag according to the following chemical reaction? 3Ag(s) + Al(NO3)3 → 3AgNO3 + Al(s)