Gases, Solutions, Acids, and Bases Study Guide

Unit B: Forms of Matter: Gases

How do observations of gases relate to scientific models?
What is the relationship among pressure, temperature, volume, and amount of a gas?
How is the behavior of gases used in various technologies?

4.1 Empirical Properties of Gases

Importance of gases and technologies relying on their properties.
Kinetic Molecular Theory (KMT): Explains gas properties at ordinary temperatures and pressures.
Motion of molecules differs in solid, liquid, and gas states.
- Solid
- Liquid
- Gas
According to KMT:
- Gas molecules are in constant motion.
- The distance between gas molecules is large.
- The kinetic energy of molecules depends on temperature.
Properties distinguishing gases:
- Gases are compressible.
- Gases expand as temperature increases.
- Gases diffuse easily.
- Gases have lower densities than solids and liquids.
- Gases mix evenly and completely; considered homogeneous.

Measuring Gases: Volume and Pressure

Volume: Amount of space a substance occupies.
- Measured in mL (convenience) or L (SI units).
- $1 \text{ cm}^3 = 1 \text{ mL}$
Pressure: Force per unit area exerted by moving particles colliding with container walls.
- SI unit: pascal (Pa).
- Gases often measured in kilopascals (kPa).
- Other units: atmospheres (atm), millimeters of mercury (mmHg), Torricelli (torr), bar, pounds per square inch (PSI).
Unit Conversions:
- Pascal (Pa): $1 \text{ Pa} = 1 \text{ N/m}^2$
- Atmosphere (atm): $1 \text{ atm} = 101 \text{ kPa}$
- Millimeters of mercury (mmHg): $760 \text{ mmHg} = 1 \text{ atm}$
- Torricelli (torr): $1 \text{ torr} = 1 \text{ mmHg}$
- Pounds per square inch (PSI): $1 \text{ PSI} = 6895 \text{ Pa}$
- Bar (bar): $1 \text{ bar} = 100 \text{ kPa}$
Conversion examples provided.

SI Prefixes

Prefixes adjust units for convenient measurement.
Prefix Table:
- Tera (T): $10^{12}$
- Giga (G): $10^9$
- Mega (M): $10^6$
- Kilo (k): $10^3$
- Milli (m): $10^{-3}$
- Micro ($\mu$): $10^{-6}$
- Nano (n): $10^{-9}$
- Pico (p): $10^{-12}$
Conversion examples provided.

Standard Conditions for Gases

Standard Temperature and Pressure (STP): $0 \text{ }^\circ\text{C}$ and 101.325 \tet{ kPa}
Standard Ambient Temperature and Pressure (SATP): $25 \text{ }^\circ\text{C}$ and $100 \text{ kPa}$

The Relationship between Pressure and Volume: Boyle’s Law

Boyle's Law: As pressure on a gas increases, volume decreases proportionately if temperature and amount of gas remain constant.
- Volume of a gas is inversely proportional to its pressure.
- $P1V1 = P2V2$
 - 1 = initial conditions
 - 2 = final conditions
Examples provided for Boyle's Law calculations.

The Relationship between Temperature and Volume: Charles’ Law

Charles' Law: As the temperature of a gas increases, the volume increases proportionately if pressure and amount of gas remain constant.
- The temperature of a gas is directly proportional to the volume of a gas.
- $\frac{V1}{T1} = \frac{V2}{T2}$
Temperature is a measure of average kinetic energy.
Absolute zero: Point of zero kinetic energy, theoretically -273°C.
Kelvin temperature scale: No degree symbol is used for Kelvin.
- Celsius to Kelvin conversion: $K = \text{ }^\circ\text{C} + 273$
- Kelvin to Celsius conversion: $\text{ }^\circ\text{C} = K - 273$
Examples provided for Charles' Law calculations.

The Relationship between Temperature, Volume, and Pressure: The Combined Gas Law

Combined Gas Law: For a fixed amount of gas there is a relationship between volume, temperature, and pressure:
- $\frac{P1V1}{T1} = \frac{P2V2}{T2}$
Examples provided for Combined Gas Law calculations.

4.2 Explaining the Properties of Gases

Early gas studies were empirical.
Kinetic molecular theory explains gas properties.
Why are gases compressible?
- Large spaces between gas molecules allow them to be crowded together.
What is gas pressure?
- The sum of all the forces exerted by the gas molecules when they collide with the walls of the container.
Use KMT to explain Boyle’s Law:
- As volume decreases, molecules collide with walls more frequently, increasing pressure.
Use KMT to explain Charles’ Law:
- As temperature increases, molecules move faster and occupy more space, increasing volume.
KMT explains physical but not chemical properties.
Avogadro’s Theory: Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.

4.3 Molar Volume of Gases

Mole: measurement of the number of particles in a sample.
- $1 \text{ mole} = \text{Avogadro’s number} = 6.02 \times 10^{23} \text{ atoms or molecules}$
- $1 \text{ mol of He(g)} = 6.02 \times 10^{23} \text{ atoms}$
- $1 \text{ mol of CO}_2\text{(g)} = 6.02 \times 10^{23} \text{ molecules}$
Guy-Lussac’s Law and Avogadro’s Theory can be combined.
Molar volume (Vm): the volume that one mole of a gas occupies; is the same for all gases at the same temp and pressure.
- Vm at STP = 22.4 L/mol
- Vm at SATP = 24.8 L/mol
Molar volume can calculate the chemical amount of a gas.
- $n = \frac{V}{V_m}$
  - n: chemical amount (moles)
  - V: volume of gas (L)
  - Vm: molar volume (L/mol)
Examples provided for molar volume calculations.

Molar Volume and Molar Mass

Relate mass to chemical amount using molar mass.
- $n = \frac{m}{M}$
  - n: chemical amount (moles)
  - m: mass (g)
  - M: molar mass (g/mol)
Molar Mass (M): mass of one mole of a sample (g/mol).
- Chemical formula and number of atoms are needed.
- Multiply the number of each type of atom by the atomic mass, and add together.
Examples provided for molar volume and molar mass calculations.

The Law of Combining Volumes

Gay-Lussac’s Law (Law of Combining Volumes): Volumes of gases of reactants and products in chemical reactions are always in ratios of whole numbers if measured at the same temperature and pressure.
Not all products/reactants need to be gases.
Reacting volumes of gases are in whole number ratios, like coefficients in a balanced equation.
Correct balanced reaction needed!
Examples provided for Law of Combining Volumes.

4.4 The Ideal Gas Law

All empirical properties of gases assumed to apply perfectly to all gases: ideal gas.
Ideal gas: a hypothetical gas that obeys all gas laws perfectly under all conditions.
- The size of an ideal gas is negligible.
- There are no forces of attraction between ideal gas molecules.
- Ideal gases do not condense when cooled.
- Ideal gas molecules undergo perfectly elastic collisions in which no energy is lost.
Real gas: an actual gas that condenses when cooled and deviates from the gas laws under certain conditions.
- High pressure (> 1 MPa): Intermolecular attraction leads to condensation.
- Low temperature: Intermolecular attraction leads to condensation.
Real gases behave like ideal gases at relatively low temperatures and pressures (e.g., STP and SATP).

The Ideal Gas Law Equation

Ideal-gas equation describes the interrelationship of pressure, temperature, volume, and chemical amount (moles).
- Boyle’s law: $V \propto \frac{1}{P}$
- Charles’ law: $V \propto T$
- Avogadro’s theory: $V \propto n$
$V \propto \frac{nT}{P}$
$PV = nRT$
- P: pressure (kPa)
- V: volume (L)
- n: chemical amount (mol)
- R: ideal gas constant = $8.314 \frac{L \cdot kPa}{mol \cdot K}$
- T: temperature (K)
Examples provided for Ideal Gas Law

Unit C: Matter as Solutions, Acids & Bases

Covers topics: mixtures, dissolving, concentrations, dilutions, ionization/dissociation, solubility, equilibrium, acid/base nomenclature, pH, Arrhenius definitions, neutralization, strong/weak acids/bases, and titrations.

Chapter 5.1: Solutions and Mixtures

Solution: a homogeneous mixture with uniform composition where the components are not visible.
- Composed of solute and solvent.
  - Solute: substance being dissolved.
  - Solvent: substance doing the dissolving.
Solutions are not only aqueous; they can have solutes and solvents that are gases, liquids, or solids.
- Gasoline: hydrocarbons
- Air: oxygen in nitrogen
- Bronze alloy: tin in copper
We will focus on aqueous solutions.

Aqueous Solutions

Aqueous Solution: Any solution with water as the solvent; denoted with the (aq) symbol after the chemical formula.
Most chemical reactions occur in a water environment, therefore we need to understand physical and chemical properties of aqueous solutions
Why are chemicals in solution?
- Easier to handle chemicals
- Easier to complete reactions
- Easier to control reactions

Properties of Aqueous Solutions: Solubility

Solubility: amount of solute that can be dissolved in a quantity of solvent.
- Very soluble = HIGH Solubility = (aq)
- Not very soluble = LOW Solubility = (s)
- Solubility of a solute is the concentration of solute that dissolves in a given quantity of solvent at a given temperature.
Every substance has its own unique solubility data = grams per 100 mL of water
- Insoluble: less than 0.1 g/100 mL
- Slightly Soluble: between 0.1 and 1.0 g/100 mL
- Soluble: greater than 1.0 g/100 mL
Example: NaCl is 36 g/100 mL
Saturated Solution: contains the maximum amount of dissolved solute at a given temperature; undissolved particles are in the solution.
Unsaturated Solution: does not have the maximum amount of solute in it.
Range of Solubility: degree of strength based on attraction between solute particles and solvent.

Solubility in Water

ELEMENTS generally have low solubility.
IONIC COMPOUNDS solubility can be predicted from a solubility chart.
- Locate anion at the top; decide which row the cation is located in
- Top row = high solubility (aq); Bottom row = low solubility (s)
Trends on the Solubility Chart:
- All group one ions
- All compounds containing: $NH4^+, NO3^-, ClO3^-, ClO4^-$
MOLECULAR COMPOUNDS solubility can be predicted based on polarity.
- “Like dissolves like” so for a substance to dissolve in water, which is polar, the substance must also be polar.

Conductivity

Electrolyte: an aqueous solution that conducts electricity:
- highly soluble ionic compounds
- soluble ionic hydroxides
- acids
NON-electrolyte: an aqueous solution that does not conduct electricity:
- molecular compounds and elements
Predict the ionic or molecular and electrolyte/non-electrolyte for each substance
Diagnostic Test: tested with simple conductivity meter.

Indicators

Acids and bases distinguished by their empirical properties in aqueous solutions.
Acidic Solutions: blue litmus turns red; pH <7; conductive
- Hydrogen ion (proton) containing compounds:
Basic Solutions: red litmus turns blue; pH >7; conductive
- Ionic hydroxides:
Neutral Solutions: no change in colour of litmus paper, pH =7
- Ionic; conductive
- Molecular; not conductive
Diagnostic Test: tested using red/blue litmus paper or a pH meter.

Qualitative Chemical Analysis

Use known diagnostic tests to determine the identity of four unknown solutions
Problem: Which of the solutions (labeled A, B, C, and D) is calcium chloride, citric acid, glucose, and calcium hydroxide?

Qualitative Analysis by Colour

Qualitative Analysis by Colour: Color of a solution or a flame produced identifies ions present.
NOTE: ion/flame color for various substances can be found on the back of your Periodic Table.
Examples of using color to identify substances

Chapter 5.2: Explaining Solutions

Arrhenius’ Theory: when a substance dissolves, particles separate and disperse.
- Determines the type of dissolving particles (electrolyte, nonelectrolyte)
- Electrolyte: releases charged particles; ionic compounds, acids, and bases
- Non-electrolyte: disperses w/o neutral particles; molecular compounds
Aqueous reactions require knowing major entities present when any substance is in water; use dissociation.
Dissociation: separation of ions when an ionic compound dissolves in water.

Arrhenius’s Theory

Arrhenius’s theory and how it applies to acids and bases: Acids are substances that ionize in aqueous solution to form hydrogen ions, and bases are substances that dissociate to form hydroxide ions in aqueous solution.
Acids ionize in water to produce hydrogen ions ○ Acid → $H^+ (aq) + anion$
Bases dissociate in water to produce hydroxide ions ○ Base → $Cation + OH^- (aq)$
Acid properties related to $H^+$ and basic properties related to $OH^-$
Tiny proton exist on its own forming a hydronium ion, $H_3O^+ (aq)$
Limitations of the Arrhenius Theory of Acids and Bases: does not explain certain substances failing to produce neutral solutions!

Modified Arrhenius Theory

Modified Arrhenius Theory: involves the collision of dissolved substances with water molecules.
Remember that all acidic and basic substances are aqueous solutions, so the particles will constantly be colliding with, and may also react with, water molecules.
An acid reacts with water to produce $H3O^+ (aq)$ in aqueous solution $Acid + H2O → H_3O^+ (aq)$
A base dissociates and reacts with water to produce $OH^- (aq)$ in aqueous solution $Base + H_2O → OH^-(aq)$

Acid Nomenclature

Two systems for naming acids:
- IUPAC Nomenclature: names the acid as though it were an ionic compound and adds aqueous in front of it.
- Classical System:
  - Hydrogen ide = hydroic acid
  - Hydrogen ate = ___ic acid
  - Hydrogen ite = ___ous acid
Practice using the modified Arrhenius theory

Strength of Acids and Bases

Acids with the same initial concentration can have different degrees of acidic properties!
Strong and Weak Acids:
- Strong Acid Reacts completely with water; ex: $HCl + H2O → Cl^- + H3O^+</li><li>Weak Acid Reacts incompletely with water H3COOH+ H2O ↔ H3O+ + CH3COOH-Many [H3O+ (aq)] ions in solution / Few [H3O+ (aq)] ions in solution; High conductivity / Low conductivity; Relatively low pH / Relatively high pH (closer to 7) High rate of reaction with active metals/carbonates / Lower rate of reaction with active metals/carbonates Few strong acids / Many weak acids - all the ones that aren’t strong!</li></ul></li><li>Strong and Weak Bases:<ul><li>Strong Base Dissociates completely to release OH- (aq) ions / Weak Base Reacts incompletely with water to produce few OH- ionsMOH(aq)→ M+ (aq) + OH- (aq) Where M = metal ion / B(aq) + H2O(l) ↔ OH- (aq) + Balancing Entity Where B = base (molecule/ ion)</li></ul></li><li>Polyprotic Substances:<ul><li>Monoprotic Acids: 1 acidic hydrogen Ex: HCl</li><li>Polyprotic Acids: acids w more than one acidic hydrogen Ex: H3PO4</li></ul></li><li>Monoprotic Bases vs Polyprotic Bases<ul><li>Barium examples</li></ul></li></ul><h3 id="5b733d4b-808a-4ae4-9dc0-a0a852cf020a" data-toc-id="5b733d4b-808a-4ae4-9dc0-a0a852cf020a" collapsed="false" seolevelmigrated="true">Dissociation Equations</h3><ul><li>Dissociation Equations for soluble ionic compounds: cation(aq), anion(aq), H2O(l)<ul><li>Sodium fluoride is placed in water Dissociation Equations NaF (s) → Na+ (aq) + F- (aq) ; Major Entities Na+(aq), F-(aq), H2O(l)</li><li>Insoluble ionic compounds: Major Entities: compound(s), H2O(l) Dissociation Equations MgCO3 (s) → MgCO3 (s); Major Entities: MgCO3 (s), H2O (l)</li><li>Molecular Compounds: *do NOT dissociate in water; Major Entities: H2O(l) and molecule state will depend on polarity</li><li>Acidic Compounds: Arrhenius’ theory states that acids ionize in water to produce H+ (aq) ions and negative ions; Modified Arrhenius’ theory states that acids react with water to produce H3O+ (aq); Both are acceptable when writing the dissociation equation for acids, however we typically write dissociation equations based on Arrhenius’ original theory</li></ul>Strong Acids: HIGH ionization, Major Entities: H+ (aq), anion, H2O(l)HCl (aq) → H+ (aq) + Cl- (aq) ; Major Entities H+ (aq), Cl- (aq), H2O(l)Weak Acids: low degree of ionization and are weak conductors; Major Entities: acid(aq), H2O(l)CH3COOH (aq) ↔ H+ (aq) + CH3COO- (aq); Major Entities: CH3COOH (aq), H+ (aq), CH3COO- (aq),H2O(l)</li></ul>NOTE: Summary of Dissociation Equations and Major Entities<h3 id="e9ac51e7-4eb3-406d-a601-a9a31fb50c0d" data-toc-id="e9ac51e7-4eb3-406d-a601-a9a31fb50c0d" collapsed="false" seolevelmigrated="true">Energy Changes in Dissolving</h3><ul><li>Difficulty predicting whether the dissolving of a solute will be endothermic or exothermic; We only observe the net energy change. : endothermic / exothermic</li><li>Breaking bonds requires energy endothermic; : Forming new bonds releases energy exothermic;</li><li>The overall energy in the dissolving process requires considering both energy absorbed and released.</li><li>Generalizations for Solubility in Water:</li></ul><h3 id="4aaead12-0d7b-46df-a77e-ffc55ab0fc1b" data-toc-id="4aaead12-0d7b-46df-a77e-ffc55ab0fc1b" collapsed="false" seolevelmigrated="true">Generalization</h3><ul><li>Solids: energy is needed to break bonds holding the solute together therefore as temperature increases; the solubility of a solid also increases. : Gases: at high temperatures a dissolved gas gains energy and escapes from the liquid (solvent); therefore the solubility of a gas decreases at higher temperatures; Gases have decreases solubility at increases temperatures and increases pressures : Liquids: Difficult to generalize effect of temp: 1) polar liquids have increase solubility at increase temperature and are said to be miscible. 2) nonpolar liquids do not dissolve and form a separate layer and are said to be immiscible in water. : Elements: generally have low solubility in water</li></ul><h3 id="7619c166-f158-47a6-94bd-44e84cb637cb" data-toc-id="7619c166-f158-47a6-94bd-44e84cb637cb" collapsed="false" seolevelmigrated="true">Equilibrium</h3><ul><li>DynamicEquilibrium: Amount of undissolved solute at the bottom remains unchanged; undissolved dissolves and dissolved precipitates out of solution and crystallize and is said to be in in a state of dynamic equilibrium. : Equilibrium occurs when a process (dissolving) and the reverse process (crystallization) take place at the same rate in a closed system; Representing Dissolving, Crystallization, and Equilibrium for a Saturated Solution: : Dissolving: X(s) → X(aq) : Crystallization: X(aq) → X(s) : Equilibrium: X(s) ↔ X(aq)</li></ul><h3 id="a9931ba2-aa3a-4973-be05-d874db719db5" data-toc-id="a9931ba2-aa3a-4973-be05-d874db719db5" collapsed="false" seolevelmigrated="true">Chapter 5.3: Solution Concentration/ a way to compare amt</h3><ul><li>Solution contains solute+ solvent 1) solute: substance particle. 2) Solvent dissolves solute : Concentration (quantity;Typically, concentration = quantity of solute_ quantity of solution) 1)Dilute if less particles per unit volume. 2)Concentrated if more particles per unit volume.</li></ul><h3 id="bf96949f-05f2-4d72-a182-c676438c75b5" data-toc-id="bf96949f-05f2-4d72-a182-c676438c75b5" collapsed="false" seolevelmigrated="true">Three Main Methods of Concentration</h3><ul><li>1.AmountConcentration/Molarity/MolarConcentration: the chemical amount (moles) of solute dissolved in one litre of solution.$ FORMULA: 𝑐 = 𝑛/𝑣 $Units: mol/L Where: c = concentration (mol/L); n = number of moles (mol); V = volume (L)</li><li>2. Parts per million (ppm):used for dilute solutions; altered ppm units by altering ppm to incorporate a factor of FORMULA:$ 𝑐=\frac{𝑚𝑠𝑜𝑙𝑢𝑡𝑒(𝑚𝑔)}{𝑚𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛(𝑘𝑔)} $Units: mg/L, mg/kg, ppm, (1 ppm = 1 mg/L = 1 mg/kg) Because very dilute aqueous solutions are similar to pure water, their densities are considered to be the same (1 g/mL), therefore:$ 1 \text{ ppm} = 1 \frac{g}{10^6 mL} = 1 \frac{mg}{L} = 1 \frac{mg}{kg}$$\n 3. percent concentration= solute / solution X 100%: 1) percent volume by volume concentration (% v/v):volume in mL that dissolves in every 100 mL of solution. 2) percent by mass/ volume(% w/v= mass solute=grams that dissolves in every 100 mL of solution.3) percent by mass/mass (% w/w) mass solute=grams dissolved in every 100 grams of solution.
Concentration of Ions
- Solutions of ionic compounds and strong acids are able to conduct electricity therefore there must be ions present in solution * Molar concentration of ions in solution depends on the number of ions making up the compound. Examples: Na3PO4 (aq) / CaCl2 (aq)
 Follow steps for calculating Ion Concentration: １) Write a balanced dissociation/ionization equation. ２) The ion concentration is always a whole number multiple of the compound concentration. NOTE: Square brackets around an [ion] or formula indicate the concentration of the substance within the brackets.
Ions in Solution
- Problems to practice
Chapter 5.4: Preparation of Solutions
- When a solid, you must calculate the amount of mass required to form the desired concentration and then mix that amount carefully with water following the steps outlined:
 : Standard Solution: solution w known concentration. Mass solid reqd / Mass solute is in a clean/dry beaker/Dissolved solid in distilled water/Transfer solution to a clean flask/ Add distilled to bottom of meniscus/ Invert and mix
Calculating the mass needed can be done using the following equations that we know: 1)C = n/V 2) m = Mn / Example: What mass of lead (II) nitrate must be used to produce 350 mL of solution with concentration 0.125 mol/L?
Methods of Concentration:
- Dilution = (decreasing solution concentration by adding solution) / During the process of dilution, the volume increases while the concentration decreases.
 1) Calculating mass needed: mass solution reqd 2) Add distilled water that is half way reqd. 3) Reqd volume of solution using a pipette . 4) Add stock solution to the clean volumetric flask. 5)Solution of meniscus must be on line on a flask 6) Invert and mix / C1V1 = C2V2 / This works because the quantity of solute is not changing ni = nf *On tests, you must not only calculate the solution needed, but needed procedures/ when diluting all concentrated reagents, especially acids, ALWAYS add the stock solution TO the quantity of water… not the other way around!
PH
- calculating power of hydrogen w 2 methods! 1 trace amt of hydronium hydrozide ion from is water that does ionize to small degree/ hydrate protin/1) H20= hydronium+ hyroxide = acid + base) 2 water molecules + hydronium= concentration 1 ^-7 power
 acid increase-concentration ; increase hydronium/ increase acid // hydrozide in base solution ;/ increase conc base PH increase/ decrease concentration power hydrogen is shortmethod communication hydrogen concentration powerhydrogen = hyrdounium . in power hydrogen is 0 -14scale / hydronium concentration is neutral-concentration .
PH Calculations
- PH is defined as power to hydronium and the hydronium concentration / Example: GIVEN concentration CALCULATE PH = POWER HYDROGEN CONCENTRATION!/ PH has not unit the PH is negative on log scale of hdyronium = power + /significant fig rules are normal and concentration sig figs sig figs / the number sig fig that comes after decimal PH and ph needs to follow sig fig /sample; questions and chart!
 : Calculating pH and pOH. Using the formulas, pH(H3O+ (hydronium ion))=-log (H30). pOH ( hydroxide)=- log.1)pOH +PH =14!!!!/2) 14+log (H3O)=POH!!!/3) OH 14+ log ( OH ) =PH!!!
- Calculating phOH/ acid base indicator: is substance changer color is related ph change. Indicator exists in 2 colors 1 of color and another / Indicators PH range inside cover

Gases, Solutions, Acids, and Bases Study Guide

Unit B: Forms of Matter: Gases

4.1 Empirical Properties of Gases

Measuring Gases: Volume and Pressure

SI Prefixes

Standard Conditions for Gases

The Relationship between Pressure and Volume: Boyle’s Law

The Relationship between Temperature and Volume: Charles’ Law

The Relationship between Temperature, Volume, and Pressure: The Combined Gas Law

4.2 Explaining the Properties of Gases

4.3 Molar Volume of Gases

Molar Volume and Molar Mass

The Law of Combining Volumes

4.4 The Ideal Gas Law

The Ideal Gas Law Equation

Unit C: Matter as Solutions, Acids & Bases

Chapter 5.1: Solutions and Mixtures

Aqueous Solutions

Properties of Aqueous Solutions: Solubility

Solubility in Water

Conductivity

Indicators

Qualitative Chemical Analysis

Qualitative Analysis by Colour

Chapter 5.2: Explaining Solutions

Arrhenius’s Theory

Modified Arrhenius Theory

Acid Nomenclature

Strength of Acids and Bases

Concentration of Ions

Ions in Solution

Chapter 5.4: Preparation of Solutions

Methods of Concentration:

PH

PH Calculations