Ch1000 Fundamentals Of Chemistrymodule 4 Chapter 15arrhenius Aci ✓ Solved

CH1000 Fundament als of Chemistry Module 4 – Chapter 15 Arrhenius Acids Arrhenius Acid: An acid solution contains an excess of H+ ions. Common Properties of Acids 1. Sour taste 2. Turns litmus paper pink 3. Reacts with: Metals to produce H2 gas Bases to yield water and a salt Carbonates to give carbon dioxide Arrhenius Bases Arrhenius Bases: A basic solution contains an excess of OH– ions.

Common Properties of Bases 1. Bitter/caustic taste 2. Turns litmus paper blue 3. Slippery, soapy texture 4. Neutralizes acids Brà¸nsted-Lowry Acids and Bases Lewis Acid-Bases Summary of the Acid/Base Theories Reactions of Acids Base Reactions Bases can be amphoteric (act as either Brà¶nsted acids or bases) In general: Zn(OH)2 (aq) + 2 HBr (aq) ZnBr2 (aq) + 2 H2O (l) As a base: NaOH and KOH can also react with metals.

2 NaOH (aq) + 2 Al (s) + 6 H2O (l) 2 NaAl(OH)4 (aq) + 3 H2 (g) base + metal + water salt + hydrogen Zn(OH)2 (aq) + 2 NaOH (aq) Na2Zn(OH)4 (aq) As an acid: Salts Salts: products from acid-base reactions. HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l) Salts are ionic compounds. Salts contain a cation (a metal or ammonium ion) derived from the base and an anion (excluding oxide or hydroxide ions) derived from the acid. Salts are generally crystalline compounds with high melting and boiling points. Electrolytes and Nonelectrolytes Electrolytes: compounds that conduct electricity when dissolved in water.

Nonelectrolytes: substances that do not conduct electricity when dissolved in water. Ion movement causes conduction of electricity in water. 3 classes of compounds, acids, bases, and salts are electrolytes because they produce ions in water when they dissolve. Comparing Solution Conductivity (Sugar solution) (Salt solution)(Distilled water) Dissociation of Electrolytes Salts dissociate into their respective cations and anions when dissolved in water. Hydrated sodium (purple) and chloride (green) ions The negative end of the water dipole is attracted to the positive Na+ ion.

When NaCl dissolves in water, each ion is surrounded by several water molecules. The permanent dipoles in the water molecules cause specific alignment around the ions. NaCl (s) Na + (aq) + Cl- (aq) Electrolyte Ionization Ionization: process of ion formation in solution. Ionization results from the chemical reaction between a compound and water. Acids ionize in water, producing the hydronium ion (H3O+) and a counter anion.

Bases ionize in water, producing the hydroxide ion (OH-) and a counter cation. HCl (g) + H2O (l) H3O + (aq) + Cl- (aq) H3PO4 (aq) + H2O (l) H2PO4 - (aq) + H3O + (aq) NH3 (aq) + H2O (l) OH - (aq) + NH4 + (aq) Strong and Weak Electrolytes Strong electrolytes: undergo complete ionization in water. Example: HCl (strong acid) Weak electrolytes: undergo incomplete ionization in water. Example: CH3COOH (weak acid) HCl (left) is 100% ionized. CH3COOH exists primarily in the unionized form.

HF (aq) + H2O (l) F - (aq) + H3O + (aq) Double arrows indicate incomplete ionization (typically weak electrolytes). Salts Salts can dissociate into more than 2 ions, depending upon the compound. A 1 M solution of NaCl produces a total of 2 M of ions. A 1 M solution of CaCl2 produces a total of 3 M of ions. NaCl (s) Na+ (aq) + Cl- (aq) 1M 1M 1M CaCl2 (s) Ca 2+ (aq) + 2 Cl- (aq) 1M 1M 2M Colligative Properties of Electrolyte Solutions Colligative properties: depend only on the number of moles of dissolved particles present.

This must be taken into consideration when calculating freezing point depression or boiling point elevation due to the presence of solute particles. Example: What is the boiling point elevation of a 1.5 m aqueous solution of CaCl2? (Kb for water is 0.512 ºC/m). Because CaCl2 is a strong electrolyte, 3 mol of ions (1 mol Ca2+ and 2 mol Cl- ions) will be present in the solution. ΔTb = 1.5 m CaCl2 = 2.3 ºCà— 3 mol ions 1 mol CaCl2 0.512 ºC 1 m à— Autoionization of Water Pure water auto(self) ionizes according to the reaction: Based on the reaction stoichiometry: Concentration H3O+ = Concentration OH– = 1 x 10–7 M [H3O+] x [OH–] = (1 x 10–7)2 = 1 x 10–14 When acid or base is present in water, [H3O+] and [OH-] change.

In acidic solutions, [H3O+] > [OH–]. In basic solutions, [H3O+] < [OH–]. H2O (l) + H2O (l) H3O + (aq) + OH– (aq) Introduction to pH The pH scale Increasing acidity Increasing basicityHigh H3O + Low OH- Low H3O + High OH- In pure water, [H3O +] = 1 x 10-7 M, so pH = - log(1 x 10-7) = 7 pH = - log[H3O +] pH Calculations pH = - log[H3O +] [H3O +] = 1 x 10-5 M [H3O +] = 2 x 10-5 M If exactly 1 Exponent = pH pH = 5 If a number between 1 and 10 The pH is between the exponent and next lowest whole number pH = 4.7 Generalizations [H3O +] = 10-pH pH Neutralization General Reaction Example Overall Ionic Equation: H+(aq) + Cl- (aq) + Na+ (aq) + OH- (aq) Na+ (aq) + Cl- (aq) + H2O (l) All species are included; soluble compounds shown as ions.

Net Ionic Equation: Spectator ions (orange) are removed from both sides. acid + base salt + water HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l) H+ (aq) + OH- (aq) H2O (l) Titration Titration: experiment where the volume of one reagent (titrant) required to react with a measured amount of another reagent is measured. Titrations allow the amount of an acid or base present in a sample to be determined. Indicators are used to signal the endpoint of a titration, the point when enough titrant is added to react with the acid/base present. Burets deliver measured amounts of the titrant into a solution of the unknown reagent. Endpoin t Net Ionic Equations Rules for Writing Net Ionic Equations 1.

Strong electrolytes are written as the corresponding ions. Example: NaOH (aq) is written as Na+(aq) + OH-(aq) 2. Weak electrolytes and nonelectrolytes are written as molecules. Example: CH3OH(aq), CH3COOH(aq), etc. 3.

Solids and gases are written as their molecular forms. 4. The net ionic equation does not include spectator ions. 5. The net ionic equation must balance atoms and charge.

Acid Rain 1. Emission of nitrogen or sulfur oxides. 2. Transportation of these chemicals throughout the atmosphere. 3.

Chemical reaction of the oxides with water. 4. This forms sulfuric and nitric acids. 5. Precipitation carries the acids to the ground.

General Process for Acid Rain Formation: Acid rain: atmospheric precipitation more acidic than typical. Reading Review 1. What ion is present in an acid? What ion is present in a base? 2.

What products are produced in a reaction between an acid and a metal oxide? 3. What type of reaction produces salts? 4. What is the pH range for acids?

5. What is titration? Slide 1 Arrhenius Acids Arrhenius Bases Brà¸nsted-Lowry Acids and Bases Summary of the Acid/Base Theories Reactions of Acids Base Reactions Salts Electrolytes and Nonelectrolytes Dissociation of Electrolytes Electrolyte Ionization Strong and Weak Electrolytes Salts Colligative Properties of Electrolyte Solutions Autoionization of Water Introduction to pH pH Calculations pH Neutralization Titration Net Ionic Equations Acid Rain Reading Review CH1000 Fundament als of Chemistry Module 4 – Chapter 13 Liquid State of Matter • Liquids are an intermediate between gases and solids • They 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. • Liquid Review: • Close contact • Some attractive forces • Fluid shape Changes in State Vaporization Liquid to Vapor Molecules in liquid state have different kinetic energies (KEs) Those with higher KEs can overcome attractive forces between particles and escape to the gas phase Sublimation Solid to Vapor Phase change from the solid to gas phase that bypasses the liquid state Condensation Vapor to liquid Molecules in the gas phase can strike the surface of a liquid and return to the liquid phase In a closed container, an equilibrium develops between molecules evaporating and condensing Vapor Pressure • Vapor Pressure is the pressure exerted by a vapor in equilibrium with its liquid phase. • Independent of the quantity of liquid or its surface area • Increases with increasing temperature • Depends on the strength of attraction between molecules in the liquid state • Volatile liquids have very weak attractive forces between molecules.

Evaporate very rapidly at ambient temperature. Have high vapor pressures as a result Measuring Vapor Pressure of a Liquid •Measure using a barometer •Vapor from the liquid exerts a force on the Hg and pushes the column downward •The difference in height relative to vacuum provides the vapor pressure for the liquid Surface Tension •Resistance of a liquid to an increase in surface area. •Molecules on a liquid surface are strongly attracted by molecules within the liquid. •Surface tension increases with increasing attractive interactions between molecules. Capillary Action Capillary Action is the spontaneous rise of a liquid in a narrow tube Cohesive forces exist between water molecules in a liquid 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. Boiling Point • Temperature at which the vapor pressure of a liquid is equal to the external pressure above the liquid. • The normal boiling point is the boiling temperature when the vapor pressure is 1 atm. Freezing/Melting Point • The freezing/melting point is the temperature at which the solid phase of a substance is in equilibrium with its liquid phase • While both solid and liquid phases are present, the temperature remains constant. • The energy is used to change the solid to the liquid phase.

Changes of State •Heat of fusion is the energy required to change 1 g of a solid at its melting point to a liquid • The heat of fusion for water is 335 J/g. • Use the heat of fusion as a conversion factor •Heat of vaporization is the energy required to change 1 g of liquid to vapor at its normal boiling point. • The heat of vaporization for water is 2259 J/g. • Use the heat of vaporization as a conversion factor Intermolecular Forces Dipole-Dipole Attractions • In covalent molecules, due to different atoms having different electronegativities, molecules are polar • When polar molecules are put together, they will align to permit interaction between oppositely polarized portions of the molecules Hydrogen Bonding • A special type of dipole-dipole attraction • One type of strong intermolecular force/attraction between molecules • To form hydrogen bonds, a compound must have covalent bonds between hydrogen and F, O, or N.

London Dispersion Forces • Interaction between nonpolar molecules and noble gases • London forces arise from uneven, instantaneous charge distributions due to electron movement in nonpolar molecules. • 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. Hydrates •Hydrates are solids that contain water molecules as part of their crystalline structure •The formula lists the anhydrous formula of the compound followed by the number of waters present per structural unit. •Hydrates are named by placing a prefix corresponding to the number of water molecules. Followed by hydrate •Hydrates will often decompose by losing water upon heating •To calculate % water in a hydrate: • Calculate the molar mass of the compound • Calculate the %water of the compound Water: A Unique Liquid • Physical properties of water • Colorless, odorless, tasteless liquid • More dense in liquid than solid phase • 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 in the 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.

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 – 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) Reading Review What are the three changes in state? What is vapor pressure? What are the three types of intermolecular forces? What molecular shape does water have? How do you know a compound is a hydrate from its formula? Slide 1 Liquid State of Matter Changes in State Vapor Pressure Measuring Vapor Pressure of a Liquid Surface Tension Capillary Action Capillary Action in Action Boiling Point Freezing/Melting Point Changes of State Intermolecular Forces Hydrates Water: A Unique Liquid Slide 15 Reading Review

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Arrhenius Acids and Bases: A Comprehensive Overview

The study of acids and bases forms a vital part of chemistry, with various theories explaining their behavior and properties. One of the most fundamental approaches is the Arrhenius theory, which defines acids as substances that produce hydrogen ions (H⁺) in solution, and bases as substances that yield hydroxide ions (OH⁻). This essay aims to elucidate the core principles surrounding Arrhenius acids and bases, their reactions, as well as related concepts such as electrolyte behavior, pH, and neutralization.

Arrhenius Acids and Bases

An Arrhenius acid is characterized by its ability to donate H⁺ ions in aqueous solution. Common properties of acids include their sour taste, ability to turn litmus paper red, and reactivity with metals to release hydrogen gas (H₂) (Atkins & Paula, 2014). The classic example of an Arrhenius acid is hydrochloric acid (HCl), which ionizes in water as follows:
\[ \text{HCl (aq)} \rightarrow \text{H}^+ (aq) + \text{Cl}^- (aq) \]
On the other hand, an Arrhenius base is defined as a substance that liberates hydroxide ions (OH⁻) in solution. Bases typically have a bitter taste, feel slippery or soapy, and turn litmus paper blue (Benson, 2011). Sodium hydroxide (NaOH) serves as a standard illustration of an Arrhenius base:
\[ \text{NaOH (aq)} \rightarrow \text{Na}^+ (aq) + \text{OH}^- (aq) \]

Reactions of Acids and Bases

Arrhenius acids and bases participate in neutralization reactions, where they react to form salt and water. For example, the reaction of hydrochloric acid and sodium hydroxide can be expressed as:
\[ \text{HCl (aq)} + \text{NaOH (aq)} \rightarrow \text{NaCl (aq)} + \text{H}_2\text{O (l)} \]
Through this equation, we can see how the ions exchange during the neutralization process (Tobias, 2007).

Role of Salts and Electrolytes

Salts, compounds formed from the reaction of acids and bases, are critical for various chemical processes. Salts consist of cations derived from bases and anions from acids. For instance, sodium chloride (NaCl), formed from NaOH and HCl, is an ionic compound with distinct physical properties (Kooyman & Leferink, 2015).
Salts, like acids and bases, act as electrolytes in solution when dissolved in water. An electrolyte is any substance that can conduct electricity via its ion content. Acids release H⁺ ions, bases release OH⁻ ions, and salts dissociate into their respective ions in water (Dasgupta, 2012). Strong electrolytes, such as HCl, dissociate completely, while weak electrolytes, such as acetic acid (CH₃COOH), only partially dissociate (Chaplin, 2006).

Ionization and Autoprotolysis of Water

Both acids and bases can undergo ionization in water, leading to the formation of hydronium ions (H₃O⁺) for acids and hydroxide ions (OH⁻) for bases (Clegg, 2004). The general ionization of water can be expressed as:
\[ \text{H}_2\text{O (l)} + \text{H}_2\text{O (l)} \rightleftharpoons \text{H}_3\text{O}^+ (aq) + \text{OH}^- (aq) \]
This reaction illustrates the autoprotolysis of water and is critical in understanding pH, which quantifies the acidity or basicity of a solution.

Understanding pH

The pH scale, which ranges from 0 to 14, is crucial for measuring the concentration of H₃O⁺ ions in solution. A pH of 7 denotes neutrality, while values below 7 indicate acidity and values above 7 signify basicity (Stang & Hennings, 2010). The relationship between H₃O⁺ concentration and pH is given by the equation:
\[ \text{pH} = -\log[H_3O^+] \]
For instance, a solution with [H₃O⁺] equal to \(1 \times 10^{-5}\, \text{M}\) will possess a pH of 5 (Kauffman, 2000).

Neutralization and Titration

Neutralization reactions not only form water and salts but can also be quantitatively analyzed via titration. Titration is a technique whereby a known concentration of titrant is added to a solution until the reaction reaches an endpoint, indicated by a change in color of an indicator (Baker, 2017). The endpoint signifies the equivalence point where the amount of acid equals the amount of base.

Summarizing Key Acid-Base Theories

Beyond the Arrhenius definition, other theories such as Bronsted-Lowry and Lewis extend the understanding of acids and bases. The Bronsted-Lowry theory posits that acids are proton donors, while bases are proton acceptors (Harris, 2007). Lewis acids and bases are defined based on electron pair acceptance and donation, expanding the definition even further (Duncan, 2013).

Conclusion

Understanding Arrhenius acids and bases forms a foundational aspect of chemistry that drives various applications in both academia and industry. From their definitions to properties, reactions, and wider theories, they play a central role in chemical science. Overall, the capacity of these substances to facilitate reactions, conduct electricity, and exist as ions in solution is essential to grasp the fundamental principles of chemistry.

References

1. Atkins, P. W., & Paula, J. de. (2014). Physical Chemistry. Oxford University Press.
2. Baker, R. (2017). Analytical Chemistry. Wiley.
3. Benson, S. W. (2011). Chemistry: Concepts and Applications. Holt McDougal.
4. Chaplin, M. (2006). Water Structure and Behavior. Water Structure and Science.
5. Clegg, S. L. (2004). Aqueous Electrolyte Solutions. University of Lund.
6. Dasgupta, P. K. (2012). Aquatic Chemistry: Chemical equilibria and rates in natural waters. Wiley.
7. Duncan, C. (2013). Chemistry for Dummies. Wiley.
8. Harris, D. C. (2007). Quantitative Chemical Analysis. W. H. Freeman.
9. Kauffman, G. B. (2000). Electrochemistry and the Future of Chemistry. Wiley-Interscience.
10. Kooyman, P. J. & Leferink, F. (2015). Ionic Compounds: The Nature of Salts. Nature Publishing.