Chapter 13-study guide PDF

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CHAPTER 13 13.1 States of matter review Gases: contain particles that are far apartin random motion and independent of one another.  Little contact  No attractive forces  Random motion Solids: contain particles very close in space and maintain a rigid shape. Significant attractive forces exist between particles.  Close contact  Strong attractive forces  Rigid shape Liquids: intermediate between gases and solids. Contain particles close to one another but have fluidity (can assume the shape of a container). Significant attractive forces exist between particles in a liquid.  Close contact  Some attractive forces  Fluid shape 13.2 Properties of liquids Evaporation Evaporation or Vaporization: Escape of molecules from the liquid to the gas phase. Liquid Vapor Kinetic Energy: Energy associated with movement. 1. Molecules in the liquid state have different kinetic energies (KE’s). 2. Those with higher KE’s can overcome attractive forces between particles and escape to the gas phase. 3. We can increase the kinetic energy (or movement of the molecules in liquid state by adding heat (boiling). Sublimation Phase change from the solid to gas phase that bypasses the liquid state. Solid vapor Examples: Fusion Molecules can go from the liquid to solid phase. This process is called “fusion” Ex. Liquid solid Under certain circumstances gas can transform directly into a solid. This process is called “deposition”. Example: Water vapor to ice - Water vapor transforms directly into ice without becoming a liquid a process that often occurs on windows during the winter months. Condensation Molecules from the liquid phase can escape to the vapor phase through evaporation.

Molecules in the gas phase can strike the surface of the liquid and return to the liquid phase. This process is called condensation. Ex: Rain: H2O(g) molecules condense into H2O(l) molecules Evaporation-Condensation Equilibrium Liquid

Vapor

In a closer container an equilibrium (or balance) develops between molecules evaporating and condensing. Vapor pressure Vapor pressure: pressure exerted by a vapor in equilibrium with its liquid phase.  Evaporation  Condensation begins  Rate of evaporation = rate of condensation Vapor pressure Vapor pressure: pressure exerted by a vapor in equilibrium with its liquid phase. 1. Independent of the quantity of liquid or its surface area. 2. Increases with increasing temperature. 3. Depends on the strength of attraction between molecules in the liquid state. (This means it depends on what types of liquids since different types of liquids will have different molecules.) Volatile liquids: very weak attractive forces between molecules. Evaporate very rapidly at ambient temperature. Have high vapor pressures as a result. Ex: Ethanol acetone methanol gasoline etc. Surface Tension Surface Tension: Resistance of a liquid to an increase in surface area. 1. Molecules on a liquid surface are strongly attracted by molecules within the liquid. 2. Surface tension increases with increasing attractive interactions between molecules. (In general this means liquids with more polar molecules will have more surface tension. Ex. Water (a polar compound) has a high surface tension. Ex. Nonpolar compounds include fats oil and petrol/gasoline have low surface tension. Capillary Action Capillary action: Spontaneous rise of a liquid in a narrow tube. Two forces are acting against each other: 1. Cohesive forces exist between water molecules in a liquid. 2. Adhesive forces exist between water molecules and the walls of the container. When the cohesive forces between molecules are less than the adhesive forces between liquid and container the liquid will move up the walls of the container. Capillary Action in Action

Shape of the meniscus reflects the relative strength of cohesive forces within the liquid and adhesive forces between the liquid and the tube. If convex:adhesive forces < cohesive forces If concave:adhesive forces > cohesive forces Boiling Point Boiling point: Temperature at which the vapor pressure of a liquid is equal to the external pressure above the liquid. “Normal” boiling point: The temperature when the vapor pressure is 1 atm. Normal boiling points:  Water: 100 o C  Ether: 35 o C  Ethyl Alcohol: 78 o C Vapor Pressure Curve

Freezing Point or Melting Point Freezing/melting point: the temperature at which the solid phase of a substance is in equilibriumwith its liquid phase. For H2O (at 1 atm) it is 0°C 273 K or 32°F solid liquid While both phases are present the temperature will not change even if you are adding more heat. The energy (heat) is being used to change the solid to the liquid phase. After the ice/water mixture becomes liquid then the temperature will start to rise again. 13.4 Changes of State Energy and Phase changes 1. Heat of fusion: energy required to change 1 g of a solid at its melting point to a liquid. Heat of melting: (Exact opposite of heat of fusion). 2. Heat of vaporization: energy required to change 1 g of a liquid to vapor at its normal boiling point. Heat of Condensation: (Exact opposite of heat of fusion). Solid Liquid Liquid Gas Heat of Fusion Heat of Vaporization PRACTICE The heat of fusion for water is 335 J/g. Calculate the amount of heat needed to melt 25.0 g of water.

Use the heat of fusion as a conversion factor! The heat of vaporization for water is 2259 J/g. Calculate the amount of heat needed to vaporize 25.0 g of water at 100 oC.

Changes of State 1. Heat of fusion: energy required to change 1 g of a solid at its melting point to a liquid. 2. Heat of vaporization: energy required to change 1 g of a liquid to vapor at its normal boiling point.

1. A-B: A solid is heating up. Temperature is increasing. 2. B-C: Heat of fusion. The heat (energy) is being used to change the solid to a liquid. No change in temperature. 3. C-D: The liquid is heating up. Temperature is rising. 4. D-E: Heat of vaporization. The heat (energy) is being used to change the liquid to a gas. No change in temperature. 5. E-F: The gas is heating up. Temperature is rising. 1. A-B: A solid is heating up. Temperature is increasing. ((Use q = mCs∆T)) 2. B-C: Heat of fusion. The heat (energy) is being used to change the solid to a liquid. No change in temperature. (Use q = m ∆H(fusion/melting)) 3. C-D: The liquid is heating up. Temperature is rising. ((Use q = mCl∆T)) 4. D-E: Heat of vaporization. The heat (energy) is being used to change the liquid to a gas. No change in temperature. (Use q = m ∆H(vap)) 5. E-F: The gas is heating up. Temperature is rising. (Use q = mCg∆T) Phase change practice Calculate the energy needed to convert 25.0 g of ice at 0 oC to steam at 100 oC? Given: heat of fusion = 335 J/gheat of vaporization = 2259 J/g specific heat of liquid water = 4.184 J/goC

How many joules of energy are needed to change 10.0 g of ice at 0.00 oC to water at 20.0 oC? Given: heat of fusion = 335 J/gspecific heat of liquid water = 4.184 J/goC

13.5 Intermolecular Forces Attractive forces between molecules. These forces allow for formation of liquids and solids. The degree of intermolecular forces correlates with a compound’s physical properties. Example: The stronger the interaction between molecules in a liquid the higher the boiling point (bp) and the lower the vapor pressure (vp). Intermolecular Forces 1. Dispersion Forces 2. Dipole-Dipole Forces 3. Hydrogen Bonding 4. Ion-Dipole London Dispersion Forces All molecules can interact with one another. Ex. CO2 I2 CH3CH2CH3 etc. 1. These interactions between nonpolar molecules and noble gases are called London dispersion forces. 2. London forces (also called induced dipole-dipole forces or van der Waals forces) arise from uneven instantaneous charge distributions due to electron movement in nonpolar molecules. 3. Dispersion forces are present among ALL molecules whether they are polar or nonpolar. This instantaneous dipole can then induce a dipole in a neighboring nonpolar molecule resulting in a small attraction between particles.

1. London forces are forces. 2. Generally become more important as the size of the molecule increases. Larger sizes provide more possible electrons to provide dipoles.

Which of the following molecules would have the largest London dispersion forces? Dipole-Dipole Forces Dipole-Dipole Attractions: 1. In covalent molecules due to different atoms having different electro-negativities molecules are polar. 2. When polar molecules are put together they will align to permit interaction between oppositely polarized portions of the molecules. These interactions between dipoles in different molecules are called dipole-dipole forces. The Hydrogen Bond

Water has very high melting and boiling points and heats of fusion and vaporization. These anomalous properties are due to strong attraction between water molecules due to hydrogen bondinga special type of dipole-dipole attraction. Hydrogen Bonding Hydrogen bonds: 1. A special type of dipole-dipole force that allows for strong intermolecular force/attraction between molecules. 2. But they are still are much weaker than ionic or covalent bonds which are intramolecular forces.

To form hydrogen bonds a compound must have covalent bonds between hydrogen and F O or N (very electronegative elements). 13.6 Hydrates Hydrates are solids that contain water molecules as part of their crystalline structures. The formula lists the anhydrous (without water) formula of the compound. The number of water molecules present per structural unit (water of hydration) are then

given. CaCl2∙2H2O Hydrates are named by placing a prefix corresponding to the number of water molecules followed by hydrate CaCl2∙2H2O calcium chloride dihydrate FeCl3∙6H2O iron(III) chloride hexahydrate Hydrates will often decompose by losing water upon heating. Water: A Unique Liquid Water covers 75% of the Earth’s surface; 97% of all water resides in the oceans. Water constitutes 70% of a human body by mass. Physical Properties of Water Colorless odorless tasteless liquid. More dense in liquid than solid phase (why ice floats). High boiling point high heat of fusion/vaporization due to hydrogen bonding. Structure of Water Molecules Two OH bonds are formed by the overlap of 1s orbitals on H with orbitals on the O. The molecular geometry of water is bent due to the two lone pairs on oxygen. Water has a permanent dipole due to the molecules’ shape and the polar O-H bonds. Sources of Water for Human Consumption Climate change and increased demand for fresh water make finding and sustaining sources of potable water critical for future generations. Strategies to Sustain Water Supplies 1. Reclamation of wastewater: Currently used in agriculture and industry 2. Desalination of seawater. Expensive but useful for countries near the ocean. 3. Low temperature distillation. At low pressure water’s boiling point is reduced. Less energy is required to separate the salts by boiling. 4. Combustion of H2. H2 and O2 react very exothermically to produce water. Chemistry in Action: Osmosis Osmosis: process by which water flows through a membrane from a region of more pure water to a region of less pure water. Reverse Osmosis Reverse Osmosis: process by which water flows through a membrane from a region of less pure water to a region of more pure water due to the presence of an external stimulus (typically pressure) Process often used in water purification.A semipermeable membrane is used and only water can pass through.By applying pressure only water passes through the membrane. The water is now pure! 14.1 General Properties of Solutions

Solution: a homogeneous mixture of one or more solutes and a solvent. Solute: substance being dissolved. Solvent: dissolving agent that is usually themost abundant substance in the mixture.

Properties of a True Solution 1. A homogeneous mixture of 2 or more components whose ratio can be varied. 2. The dissolved solute is molecule or ionic in size (< 1 nm). 3. Can be colored or colorless though solutions are usually transparent. 4. The solute remains dissolved and does not settle (precipitate) out of solution over time. 5. The solute can be separated from solvent by physical means (usually evaporation). 14.2 Solubility  Solubility: the amount of a substance that will dissolve in a specific amount of solvent at a given temperature. Example 27 g KBr/100g H2O at 23 oC  Miscible: when two liquids dissolve in each other.  Immiscible: when two liquids do not dissolve one another. A mixture of oil and water is immiscible.

Factor Affecting Solubility: 1) Polarity For molecular or covalent compounds it depends on the polarity of the solute and the solvent. “Like Dissolves Like” 1. Polar compounds dissolve in polar solvents. Ethanol (CH3OH) dissolves in water (HOH).

2. Nonpolar compounds dissolve in nonpolar solvents. Carbon tetrachloride (CCl4) dissolves in hexane (CH3(CH2)4CH3). Ionic Compound Solubility in Polar Solvents Majority of ionic compounds will dissolve in water due to strong ion-dipole forces. 1. The individual cations and anions are surrounded by H2O molecules (i.e. hydrated). 2. The cation is attracted to the partially negative O atom. 3. The anion is attracted to the partially positive H atoms. Factors Affecting Solubility: 2) Temperature

1. Solubility increases with temperature for most solids (red lines). 2. Solubility decreases with temperature for all gases (blue lines). As a gas increases in temperature the kinetic energy increases which means it interacts less with the liquid making it less easy to solvate. Factor Affecting Solubility: 3) Pressure Gas solubility in a liquid is proportional to the gas pressure over the liquid. (Higher pressure allows more gas molecules to dissolve in a liquid.) 

Pressure does not affect solubility of liquids or solids.

Example: A bottle of root beer is under high pressure. As the bottle opens the pressure decreases. This causes CO2 to come out of the liquid as gas bubbles. Saturate and Unsaturated Solutions There are limits to the solubility of a compound at a given temperature. Saturated solutions: contain the maximum amount of dissolved solute in a solvent. Saturated solutions are still dynamic; dissolved solute is in equilibrium with undissolved solute. undissolved solute dissolved solute 

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Unsaturated solutions: contain less than the maximum amount of possible dissolved solute in a solvent.

Supersaturated Solutions: contain more solute than needed to saturate a solution at a given temperature. How is this possible? Heating a solution can allow more to dissolve. Upon cooling to ambient temperature the solution is supersaturated. These solutions are unstable -- disturbing the solutions can cause precipitation of solute.

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Some hot packs release heat by crystallization of a supersaturated solution of sodium acetate. 14.3 Rate of Dissolving Solids 1. Effect of Particle Size:  Increasing the surface area of a solid will increase the rate of solubility.  A solid can only dissolve at a surface that is in contact with the solvent.  Since smaller crystals have a higher surface to volume area smaller crystals dissolve faster than larger ones. 2. Effect of Temperature a) Increasing the temperature normally increases the rate of dissolution of most “solid” compounds. b) Solvent molecules strike the solid surface more often causing the solid to dissolve more rapidly. c) The solute molecules are more easily separate from the solid due to a higher kinetic energy. **Increasing temperature also increases solubility of most solids. **Increasing temperature decreases solubility of gases. 3. Effect of Solute Concentration a) Rate is highest at higher concentration and decreases at lower concentration. b) As the solution approaches the saturation point the rate of solute dissolving decreases. 4. Effect of Agitation/Stirring  Stirring a solution briskly breaks up a solid into smaller pieces increasing surface area thereby increasing the rate of dissolution. Solutions: A reaction Medium The purpose of dissolving reactants in a solution is often to allow them to come in close contact to react. Example:Solid-solid reactions are generally very slow at ambient temperature KCl (s) + AgNO3 (s) No Reaction By dissolving both compounds in water the ions can collide with one another and react to form an insoluble compound.

14.4 Concentration of Solutions Qualitative Expressions of Concentrations

1. Dilute: a solution that contains a relatively small amount of dissolved solute. 2. Concentrated: a solution that contains a relatively large amount of dissolved solute. Example: A 0.1 M HCl solution is dilute acid. Example: A 12.0 M HCl solution is concentrated acid. Page 27** Mass Percent What is the mass of Na2CO3 needed to make 350.0 g of a 12.3% aqueous solution?

Mass -Volume Percent Saline is a 0.9 m/v % NaCl solution. What mass of sodium chloride is needed to make 50 mL of saline? (density = 1g/mL .)

Volume- Percent What volume of soda that is 6.0% by volume alcohol contains 200.0 mL of ethanol (CH3CH2OH)?

Molarity A common unit for solution concentration in science. Example: To prepare a 1.0 M KCl solution 1.0 mol of KCl is dissolved in enough water to make 1.0 L of solution. Molarity Practice Calculate the molarity of a solution prepared by dissolving 9.35 g of KCl in enough water to prepare a 250.0 mL solution.

How many grams of KOH are required to prepare 600.0 mL of a 0.450 M KOH solution?

Dilution Dilution: Adding a solvent to a concentrated solution to make the solution less concentrated (i.e. dilute). When a solution is diluted only the volume changes. The number of moles of solute remains constant. Moles before dilution = moles after dilution Molarity1 x Volume1 = Molarity2 x Volume2 PRACTICE What volume of 12 M HCl is needed to make 500.0 mL of a 0.10 M HCl?

14.4 Concentration Of Solutions Dilution Practice Calculate the molarity of a NaOH solution prepared by mixing 100. mL of 0.20 M NaOH with 150 mL of H2O.

Solution Stoichiometry Problem: How many mL of 0.175 M Hg(NO3)2 are needed to precipitate 2.50 g of KI? web

How many grams of AgCl will form by adding enough AgNO3 to react fully with 1500. mL of0.400 M BaCl2 solution?

15.5 Colligative Properties of Solutions Colligative Property: A solution property that depends only on the number of solute particles and not on the nature of the particles. Common Colligative Properties: 1. 2. 3. 4.

Vapor Pressure Lowering: Solutions have lower vapor pressures than pure solvent. Boiling Point Elevation: Solutions have higher boiling points than pure solvent. Freezing Point Depression: Solutions have lower freezing points than pure solvent. Osmosis and Osmotic Pressure

Vapor Pressure Lowering Dissolving a solute in a solvent lowers the vapor pressure of the solvent. As a result the solvent’s boiling point is increased (a) while the freezing point of the solvent is lowered (b).

Mol al i t y Si nc ec ol l i g a t i v ep r op e r t i e sde pe ndont hen umbe rofp a r t i c l e si nt hes o l v e n ta ndn ott he i de nt i t y , ane wc onc e n t r a t i onuni ti su s e dwhe ndi s c us s i n gc o l l i g a t i v ep r o pe r t i e s . Ex a mpl e : Wha ti st h emo l a l i t yofas ol ut i onpr e p a r e db yd i s s ol vi n g0. 10molofs t a r c hi n 0. 5 0k gofwa t e r ?

Wha t i st hemol a l i t y( m)ofas o l ut i o npr e pa r e db ydi s s ol vi n g2 . 70gofme t ha nol ( CH3 OH)i n25 . 0gofwa t e r ? Col l i ga t i v ePr o pe r t i e s Pa g e48 **** Boi l i ngPoi ntEl e v a t i o n Wha t i st heboi l i n gpoi ntofa na q ue ouss ol ut i ont ha ti s4. 00mi ns o l ut e ? Tb( p ur ewa t e r )=1 00. 0oCa ndKb=0. 512oC/ m

Fr e e z i ngPo i ntDe pr e s s i on Wha t i st hef r e e z i n gp oi ntofas ol u t i onpr e p a r e db ydi s s ol vi n g0. 10mo lofs ug a ri n0. 50 k gofwa t e r ?( Fr e e z i n gpoi ntofwa t e ri s0. 0oCa ndKf=1 . 86oC/ m)
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