# 7.5: Aqueous Solutions and Solubility: Compounds Dissolved in Water - Mathematics

Learning Objectives

• Define and give examples of electrolytes.

When some substances are dissolved in water, they undergo either a physical or a chemical change that yields ions in solution. These substances constitute an important class of compounds called electrolytes. Substances that do not yield ions when dissolved are called nonelectrolytes. If the physical or chemical process that generates the ions is essentially 100% efficient (all of the dissolved compound yields ions), then the substance is known as a strong electrolyte. If only a relatively small fraction of the dissolved substance undergoes the ion-producing process, it is called a weak electrolyte.

Substances may be identified as strong, weak, or nonelectrolytes by measuring the electrical conductance of an aqueous solution containing the substance. To conduct electricity, a substance must contain freely mobile, charged species. Most familiar is the conduction of electricity through metallic wires, in which case the mobile, charged entities are electrons. Solutions may also conduct electricity if they contain dissolved ions, with conductivity increasing as ion concentration increases. Applying a voltage to electrodes immersed in a solution permits assessment of the relative concentration of dissolved ions, either quantitatively, by measuring the electrical current flow, or qualitatively, by observing the brightness of a light bulb included in the circuit (Figure (PageIndex{1})).

Figure (PageIndex{1}): Solutions of nonelectrolytes, such as ethanol, do not contain dissolved ions and cannot conduct electricity. Solutions of electrolytes contain ions that permit the passage of electricity. The conductivity of an electrolyte solution is related to the strength of the electrolyte.

## Ionic Electrolytes

Water and other polar molecules are attracted to ions, as shown in Figure (PageIndex{2}). The electrostatic attraction between an ion and a molecule with a dipole is called an ion-dipole attraction. These attractions play an important role in the dissolution of ionic compounds in water.

Figure (PageIndex{2}): As potassium chloride (KCl) dissolves in water, the ions are hydrated. The polar water molecules are attracted by the charges on the K+ and Cl ions. Water molecules in front of and behind the ions are not shown.

When ionic compounds dissolve in water, the ions in the solid separate and disperse uniformly throughout the solution because water molecules surround and solvate the ions, reducing the strong electrostatic forces between them. This process represents a physical change known as dissociation. Under most conditions, ionic compounds will dissociate nearly completely when dissolved, and so they are classified as strong electrolytes.

Let us consider what happens at the microscopic level when we add solid KCl to water. Ion-dipole forces attract the positive (hydrogen) end of the polar water molecules to the negative chloride ions at the surface of the solid, and they attract the negative (oxygen) ends to the positive potassium ions. The water molecules penetrate between individual K+ and Cl ions and surround them, reducing the strong interionic forces that bind the ions together and letting them move off into solution as solvated ions, as Figure (PageIndex{2}) shows. The reduction of the electrostatic attraction permits the independent motion of each hydrated ion in a dilute solution, resulting in an increase in the disorder of the system, as the ions change from their fixed and ordered positions in the crystal to mobile and much more disordered states in solution. This increased disorder is responsible for the dissolution of many ionic compounds, including KCl, which dissolve with absorption of heat.

In other cases, the electrostatic attractions between the ions in a crystal are so large, or the ion-dipole attractive forces between the ions and water molecules are so weak, that the increase in disorder cannot compensate for the energy required to separate the ions, and the crystal is insoluble. Such is the case for compounds such as calcium carbonate (limestone), calcium phosphate (the inorganic component of bone), and iron oxide (rust).

## Solubility Rules

Some combinations of aqueous reactants result in the formation of a solid precipitate as a product. However, some combinations will not produce such a product. If solutions of sodium nitrate and ammonium chloride are mixed, no reaction occurs. One could write a molecular equation showing a double-replacement reaction, but both products, sodium chloride and ammonium nitrate, are soluble and would remain in the solution as ions. Every ion is a spectator ion and there is no net ionic equation at all. It is useful to be able to predict when a precipitate will occur in a reaction. To do so, you can use a set of guidelines called the solubility rules (Tables (PageIndex{1}) and (PageIndex{2})).

Table (PageIndex{1}): Solubility Rules for Soluble Substances
Soluble in WaterImportant Exceptions (Insoluble)
All Group IA and NH4+ saltsnone
All nitrates, chlorates, perchlorates and acetatesnone
All sulfatesCaSO4, BaSO4, SrSO4, PbSO4
All chlorides, bromides, and iodidesAgX, Hg2X2, PbX2 (X= Cl, Br, or I)

Table (PageIndex{2}): Solubility Rules for Sparingly Soluble Substances

Sparingly Soluble in WaterImportant Exceptions (Soluble)
All carbonates and phosphatesGroup IA and NH4+ salts
All hydroxidesGroup IA and NH4+ salts; Ba2+, Sr2+, Ca2+ sparingly soluble
All sulfidesGroup IA, IIA and NH4+ salts; MgS, CaS, BaS sparingly soluble
All oxalatesGroup IA and NH4+ salts
Special note: The following electrolytes are of only moderate solubility in water: CH3COOAg, Ag2SO4, KClO4. They will precipitate only if rather concentrated solutions are used.

As an example on how to use the solubility rules, predict if a precipitate will form when solutions of cesium bromide and lead (II) nitrate are mixed.

[ce{Cs^+} left( aq ight) + ce{Br^-} left( aq ight) + ce{Pb^{2+}} left( aq ight) + 2 ce{NO_3^-} left( aq ight) ightarrow ? onumber]

The potential precipitates from a double-replacement reaction are cesium nitrate and lead (II) bromide. According to the solubility rules table, cesium nitrate is soluble because all compounds containing the nitrate ion, as well as all compounds containing the alkali metal ions, are soluble. Most compounds containing the bromide ion are soluble, but lead (II) is an exception. Therefore, the cesium and nitrate ions are spectator ions and the lead (II) bromide is a precipitate. The balanced net ionic reaction is:

[ce{Pb^{2+}} left( aq ight) + 2 ce{Br^-} left( aq ight) ightarrow ce{PbBr_2} left( s ight) onumber ]

Example (PageIndex{1}): Solubility

Classify each compound as soluble or insoluble

1. Zn(NO3)2
2. PbBr2
3. Sr3(PO4)2

Solution

1. All nitrates are soluble in water, so Zn(NO3)2 is soluble.
2. All bromides are soluble in water, except those combined with Pb2+, so PbBr2 is insoluble.
3. All phosphates are insoluble, so Sr3(PO4)2 is insoluble.

Exercise (PageIndex{1}): Solubility

Classify each compound as soluble or insoluble.

1. Mg(OH)2
2. KBr
3. Pb(NO3)2
insoluble
soluble
soluble

## Summary

Substances that dissolve in water to yield ions are called electrolytes. Nonelectrolytes are substances that do not produce ions when dissolved in water. Solubility rules allow prediction of what products will be insoluble in water.

## Map: Introductory Chemistry (Tro)

1. The Chemical World1.1 Soda Pop Fizz
1.2 Chemicals Compose Ordinary Things
1.3 All Things Are Made of Atoms and Molecules
1.4 The Scientific Method: How Chemists Think
1.5 A Beginning Chemist: How to Succeed

2. Measurement and Problem Solving
2.1 Measuring Global Temperatures
2.2 Scientific Notation: Writing Large and Small Numbers
2.3 Significant Figures: Writing Numbers to Reflect Precision
2.4 Significant Figures in Calculations
2.5 The Basic Units of Measurement
2.6 Problem Solving and Unit Conversion
2.7 Solving Multistep Conversion Problems
2.8 Units Raised to a Power
2.9 Density
2.10 Numerical Problem-Solving Strategies and the Solution Map

3. Matter and Energy
3.2 What Is Matter?
3.3 Classifying Matter According to Its State: Solid, Liquid, and Gas
3.4 Classifying Matter According to Its Composition: Elements, Compounds, and Mixtures
3.5 Differences in Matter: Physical and Chemical Properties
3.6 Changes in Matter: Physical and Chemical Changes
3.7 Conservation of Mass: There is No New Matter
3.8 Energy
3.9 Energy and Chemical and Physical Change
3.10 Temperature: Random Motion of Molecules and Atoms
3.11 Temperature Changes: Heat Capacity
3.12 Energy and Heat Capacity Calculations

4. Atoms and Elements
4.1 Experiencing Atoms at Tiburon
4.2 Indivisible: The Atomic Theory
4.3 The Nuclear Atom
4.4 The Properties of Protons, Neutrons, and Electrons
4.5 Elements: Defined by Their Numbers of Protons
4.6 Looking for Patterns: The Periodic Law and the Periodic Table
4.7 Ions: Losing and Gaining Electrons
4.8 Isotopes: When the Number of Neutrons Varies
4.9 Atomic Mass: The Average Mass of an Element&rsquos Atoms

5. Molecules and Compounds
5.1 Sugar and Salt
5.2 Compounds Display Constant Composition
5.3 Chemical Formulas: How to Represent Compounds
5.4 A Molecular View of Elements and Compounds
5.5 Writing Formulas for Ionic Compounds
5.6 Nomenclature: Naming Compounds
5.7 Naming Ionic Compounds
5.8 Naming Molecular Compounds
5.9 Naming Acids
5.10 Nomenclature Summary
5.11 Formula Mass: The Mass of a Molecule or Formula Unit

6. Chemical Composition
6.1 How Much Sodium?
6.2 Counting Nails by the Pound
6.3 Counting Atoms by the Gram
6.4 Counting Molecules by the Gram
6.5 Chemical Formulas as Conversion Factors
6.6 Mass Percent Composition of Compounds
6.7 Mass Percent Composition from a Chemical Formula
6.8 Calculating Empirical Formulas for Compounds
6.9 Calculating Molecular Formulas for Compounds

7. Chemical Reactions
7.1 Grade School Volcanoes, Automobiles, and Laundry Detergents
7.2 Evidence of a Chemical Reaction
7.3 The Chemical Equation
7.4 How to Write Balanced Chemical Equations
7.5 Aqueous Solutions and Solubility: Compounds Dissolved in Water
7.6 Precipitation Reactions: Reactions in Aqueous Solution That Form a Solid
7.7 Writing Chemical Equations for Reactions in Solution Molecular, Complete Ionic, and Net Ionic Equations
7.8 Acid&ndashBase and Gas Evolution Reactions
7.9 Oxidation&ndashReduction Reactions
7.10 Classifying Chemical Reactions

8. Quantities in Chemical Reactions
8.1 Climate Change: Too Much Carbon Dioxide
8.2 Making Pancakes: Relationships between Ingredients
8.3 Making Molecules: Mole-to-Mole Conversions
8.4 Making Molecules: Mass-to-Mass Conversions
8.5 More Pancakes: Limiting Reactant, Theoretical Yield, and Percent Yield
8.6 Limiting Reactant, Theoretical Yield, and Percent Yield from Initial Masses of Reactants
8.7 Enthalpy: A Measure of the Heat Evolved or Absorbed in a Reaction

9. Electrons in Atoms and the Periodic Table
9.1 Blimps, Balloons, and Models of the Atom
9.3 The Electromagnetic Spectrum
9.4 The Bohr Model: Atoms with Orbits
9.5 The Quantum-Mechanical Model: Atoms with Orbitals
9.6 Quantum-Mechanical Orbitals and Electron Configurations
9.7 Electron Configurations and the Periodic Table
9.8 The Explanatory Power of the Quantum-Mechanical Model
9.9 Periodic Trends: Atomic Size, Ionization Energy, and Metallic Character

10. Chemical Bonding
10.1 Bonding Models and AIDS Drugs
10.2 Representing Valence Electrons with Dots
10.3 Lewis Structures of Ionic Compounds: Electrons Transferred
10.4 Covalent Lewis Structures: Electrons Shared
10.5 Writing Lewis Structures for Covalent Compounds
10.6 Resonance: Equivalent Lewis Structures for the Same Molecule
10.7 Predicting the Shapes of Molecules
10.8 Electronegativity and Polarity: Why Oil and Water Don&rsquot Mix

11. Gases
11.1 Extra-Long Straws
11.2 Kinetic Molecular Theory: A Model for Gases
11.3 Pressure: The Result of Constant Molecular Collisions
11.4 Boyle&rsquos Law: Pressure and Volume
11.5 Charles&rsquos Law: Volume and Temperature
11.6 The Combined Gas Law: Pressure, Volume, and Temperature
11.7 Avogadro&rsquos Law: Volume and Moles
11.8 The Ideal Gas Law: Pressure, Volume, Temperature, and Moles
11.9 Mixtures of Gases: Why Deep-Sea Divers Breathe a Mixture of Helium and Oxygen
11.10 Gases in Chemical Reactions

12. Liquids, Solids, and Intermolecular Forces
12.1 Interactions between Molecules
12.2 Properties of Liquids and Solids
12.3 Intermolecular Forces in Action: Surface Tension and Viscosity
12.4 Evaporation and Condensation
12.5 Melting, Freezing, and Sublimation
12.6 Types of Intermolecular Forces: Dispersion, Dipole&ndashDipole, Hydrogen Bonding, and Ion-Dipole
12.7 Types of Crystalline Solids: Molecular, Ionic, and Atomic
12.8 Water: A Remarkable Molecule

13. Solutions
13.1 Tragedy in Cameroon
13.2 Solutions: Homogeneous Mixtures
13.3 Solutions of Solids Dissolved in Water: How to Make Rock Candy
13.4 Solutions of Gases in Water: How Soda Pop Gets Its Fizz
13.5 Specifying Solution Concentration: Mass Percent
13.6 Specifying Solution Concentration: Molarity
13.7 Solution Dilution
13.8 Solution Stoichiometry
13.9 Freezing Point Depression and Boiling Point Elevation: Making Water Freeze Colder and Boil Hotter
13.10 Osmosis: Why Drinking Salt Water Causes Dehydration

14. Acids and Bases
14.1 Sour Patch Kids and International Spy Movies
14.2 Acids: Properties and Examples
14.3 Bases: Properties and Examples
14.4 Molecular Definitions of Acids and Bases
14.5 Reactions of Acids and Bases
14.6 Acid&ndashBase Titration: A Way to Quantify the Amount of Acid or Base in a Solution
14.7 Strong and Weak Acids and Bases
14.8 Water: Acid and Base in One
14.9 The pH and pOH Scales: Ways to Express Acidity and Basicity
14.10 Buffers: Solutions That Resist pH Change

15. Chemical Equilibrium
15.1 Life: Controlled Disequilibrium
15.2 The Rate of a Chemical Reaction
15.3 The Idea of Dynamic Chemical Equilibrium
15.4 The Equilibrium Constant: A Measure of How Far a Reaction Goes
15.5 Heterogeneous Equilibria: The Equilibrium Expression for Reactions Involving a Solid or a Liquid
15.6 Calculating and Using Equilibrium Constants
15.7 Disturbing a Reaction at Equilibrium: Le Chatelier&rsquos Principle
15.8 The Effect of a Concentration Change on Equilibrium
15.9 The Effect of a Volume Change on Equilibrium
15.10 The Effect of a Temperature Change on Equilibrium
15.11 The Solubility-Product Constant
15.12 The Path of a Reaction and the Effect of a Catalyst

16. Oxidation and Reduction
16.1 The End of the Internal Combustion Engine?
16.2 Oxidation and Reduction: Some Definitions
16.3 Oxidation States: Electron Bookkeeping
16.4 Balancing Redox Equations
16.5 The Activity Series: Predicting Spontaneous Redox Reactions
16.6 Batteries: Using Chemistry to Generate Electricity
16.7 Electrolysis: Using Electricity to Do Chemistry
16.8 Corrosion: Undesirable Redox Reactions

17.1 Diagnosing Appendicitis
17.3 Types of Radioactivity: Alpha, Beta, and Gamma Decay
17.6 Radiocarbon Dating: Using Radioactivity to Measure the Age of Fossils and Other Artifacts
17.7 The Discovery of Fission and the Atomic Bomb
17.8 Nuclear Power: Using Fission to Generate Electricity
17.9 Nuclear Fusion: The Power of the Sun
17.10 The Effects of Radiation on Life
17.11 Radioactivity in Medicine - See more at: http://www.pearsonhighered.com/educa. pzwssdaT.dpuf

This Textmap is an introductory chemistry text aimed for a single semester or quarter beginning experience to the chemistry field. This Textmap surveys some of the basic topics of chemistry and should give students enough knowledge to appreciate the impact of chemistry in everyday life and, if necessary, prepare students for additional instruction in chemistry.

## Table of Ionic Compound Solubility in Water at 25°C

Remember, solubility depends on the temperature of the water. Compounds that don't dissolve around room temperature may become more soluble in warm water. When using the table, refer to the soluble compounds first. For example, sodium carbonate is soluble because all sodium compounds are soluble, even though most carbonates are insoluble.

 Soluble Compounds Exceptions (are insoluble) Alkali metal compounds (Li + , Na + , K + , Rb + , Cs + ) ammonium ion compounds (NH 4 + Nitrates (NO 3 - ), bicarbonates (HCO 3 - ), chlorates (ClO 3 - ) Halides (Cl - , Br - , I - ) Halides of Ag + , Hg 2 2+ , Pb 2+ Sulfates (SO 4 2- ) Sulfates of Ag + , Ca 2+ , Sr 2+ , Ba 2+ , Hg 2 2+ , Pb 2+ Insoluble Compounds Exceptions (are soluble) Carbonates (CO 3 2- ), phosphates (PO 4 2- ), chromates (CrO 4 2- ), sulfides (S 2- ) Alkali metal compounds and those containing the ammonium ion Hydroxides (OH - ) Alkali metal compounds and those containing Ba 2+

As a final tip, remember solubility is not all-or-none. While some compounds completely dissolve in water and some are almost completely insoluble, many "insoluble" compounds are actually slightly soluble. If you get unexpected results in an experiment (or are looking for sources of error), remember a small amount of an insoluble compound may be participating in a chemical reaction.

## DNA Microinjection, Embryo Handling, and Germplasm Preservation

H. Greg Polites , . Carl A. Pinkert , in Transgenic Animal Technology (Third Edition) , 2014

For superovulation, we use PMSG (pregnant mare serum gonadotropin Sigma-Aldrich) and HCG (human chorionic gonadotropin, e.g., Pregnyl, Baxter Pharmaceutical Solutions LLC, Bloomington, IN) both at 5.0–7.5 units i.p. per female mouse (3–8 weeks of age) using a 26-gauge needle (e.g., B-D, 305111) attached to a 1 or 3 cc syringe. Stocks are resuspended or diluted to 25 units/mL in phosphate-buffered saline (PBS) or water, then stored at 20°C (or lower temperatures) until thawed for use. These hormones have been stored in excess of 10 years at 70°C with no apparent loss of biological activity.

With a 12- to 14-h light cycle (06:00 to 18:00 or 07:00 to 21:00 light), we administer PMSG at noon followed by at 48 h with HCG for most strains (including a variety of hybrids as the B6SJLF1, as well as outbred and inbred strains of ICR or FVB mice) or at a 26- to 48-h interval for C57BL/6 mice. Natural variation in responsiveness to administration of exogenous hormones requires that the time and dosage be titrated for any new strain in the colony, supplier of hormones, or change in the colony conditions. A biological assay related to the quality and quantity of eggs as well as mating performance (plug formation) is evaluated following hormone batch preparation.

Excessive PMSG will lead to hormone refractoriness or will increase proportions of nonfertilized, crenated, or abnormal eggs. At a proper dosage, one can reasonably expect to obtain 20–30 eggs per female mouse (dependent on age and strain, in our experience, including B6SJL, C57BL/6×DBA/2 [B6D2], and C57BL/6×C3H [B6C3H] hybrids C57BL/6 or BL/10 [and congenics], C3H or FVB inbreds and outbred Swiss [Swiss-Webster, FVB, ICR, CD-1, and ND-4]). Responsiveness of BALB/c and congenic strains developed with BALB/c mice is generally poorer with most uniform results obtained with 12-week-old donors as opposed to a greater uniformity and yield from 3- to 5-week-old donors using most other strains. Hybrid strains do show the lowest percentage of abnormal egg development, while inbred strains demonstrate higher proportions of abnormalities as a general rule. Normal eggs and some abnormal ova produced following superovulation with PMSG are illustrated in Figure 2.6C and include the following: (i) one-cell eggs with a degenerative cytoplasmic appearance (perivitelline space with multiple, fragmented polar bodies or devoid of polar bodies) (ii) highly fragmented, one-cell eggs with multiple unequal fragments in the perivitelline space (iii) precociously matured, fragmented, two-cell eggs containing unequally divided blastomeres, with cytoplasmic fragmentation and (iv) well-developed, one-cell, pronuclear-stage zygotes with a clear single (or double) polar body, two well-expanded pronuclei, and no cytoplasmic fragmentation.

## Arsenic and water: reaction mechanisms, environmental impact and health effects

 Arsenic can be found in seawater (2-4 ppb), and in rivers (0.5-2 ppb). Half of the arsenic present is bound to particles. Freshwater and seas algae contain about 1-250 ppm of arsenic, freshwater mycrophytes contain 2-1450 ppm, marine molluscs contain 1-70 ppm, marine crustaceans 0.5-69 ppm, and fishes 0.2-320 ppm (all values are based on dry mass). In some marine organisms, such as algae and shrimp, arsenic can be found in organic compounds. The legal limit for arsenic in water applied by the World Health Organization (WHO) is 10 μg/L.

In what way and in what form does arsenic react with water?

Elementary arsenic normally does not react with water in absence of air. It does not react with dry air, but when it comes in contact with moist air a layer is formed. The layer has a bronze colour, and later develops a black surface.
An example of an arsenic compounds that reacts strongly with water is orpiment. This is an amorphous arsenic compound. Reaction mechanism:

In natural water arsenic participates in oxidation and reduction reactions, coagulation and adsorption. Adsorption of arsenic to fine particles in water and precipitation with aluminium or iron hydroxides causes arsenic to enter sediments. After some time arsenic may dissolve once again consequential to reduction reactions.

Solubility of arsenic and arsenic compounds

Elementary arsenic is fairly insoluble, whereas arsenic compounds may readily dissolve. Arsenic is mainly present in watery solutions as HAsO4 2- (aq) and H2AsO4 - (aq), and most likely partially as H3AsO4 (aq), AsO4 3- (aq) or H2AsO3 - (aq).
Examples of solubility of arsenic compounds: arsenic(III)hydride 700 mg/L, arsenic(III)oxide 20 g/L, arsenic acid (H3AsO4 . 1/2 H2O) 170 g/L, and arsenic(III)sulfide 0.5 mg/L.

Why is arsenic present in water?

Arsenic compounds are abundant in the earth's crust. Particles are released during mining, and spread throughout the environment. Arsenic from weathered rocks and soils dissolves in groundwater. Arsenic concentrations in groundwater are particularly high in areas with geothermal activity. In aquatic ecosystems inorganic arsenic derived from rocks such as arsenic trioxide (As2O3), orpiment (As2S3), arsenopyrite (AsFeS) en realgar (As4S4) is most prevalent.
Arsenic is applied in different shapes and forms, and can enter water bodies as such. Large quantities of arsenic that are released from volcanic activity and from micro organisms are relatively small compared to the quantities released from for example fossil fuel combustion. Metallic arsenic is processed in lead or copper alloys, to increase hardness. The extremely toxic arsenic gas ASH3 plays an important role in microchip production. Copper arsenate (Cu3(AsO4)2.4H2O) is applied as a pesticide in viticulture, but its use is currently prohibited in many countries. Paxite (CuAs2) is an insecticide and fungicide.
Other arsenic compounds are applied as a wood preservative, in glass processing, in chemical industries, or in semiconductor technique together with gallium and indium.
Dutch painters applied arsenic as a yellow pigment. In the First World War arsenic was applied in chemical weapons. In the Vietnam War dimethyl arsenic acid was applied for the destruction of rice cultures.
Although arsenic is applied less and less, it is still present in the environment in considerable quantities. For example, near abandoned mines soil quantities of arsenic may still be up to 30 g/kg.
Arsenic was and is applied for medical purposes. In water from safe sources it probably aids curing asthma, haematological illnesses, dermatosis and psychosis. In the 19th century watery solutions of potassium arsenide (Fowler solution) were applied to treat chronic bronchial asthma and other diseases. At the beginning of the 20th century other arsenic compounds were applied to treat syphilis. Arsenic may assist in curing sleeping sickness and leukaemia.
Arsenic compounds may enter the body less specifically through food intake. This encompasses 90% of the total arsenic intake, mainly from fish products. Through fish grind in cattle feed arsenic may enter meat, and through contaminated soils it may enter plant products. In mushrooms near formed arsenic melting plants concentrations up to 50 mg/kg dry matter were found.

What are the environmental effects of arsenic in water?

Arsenic is an essential compounds for many animal species, because it plays a role in protein synthesis. It is unclear whether arsenic is a dietary mineral for humans. Arsenic toxicity is another important characteristic. The boundary concentration of arsenic is 2-46 ppm for freshwater algae. The LC50 value for Daphnia Magna is 7.4 ppm, and for the American oyster it is 7.5 ppm. These values encompass a time period of 48 hours. The chronic toxicity values for a time period of three weeks is 0.5 ppm for the large cladoceran. For rats an LC50 value of 20 mg/kg body mass was established. This is the value for the carcinogenic arsenic(III)oxide compound. This compounds also blocks enzymatic processes, increasing its toxicity. In mice, hamsters and rats the compounds was embryo toxic and teratogenic. Ferns bioaccumulate large quantities of arsenic.
Naturally, only one stable arsenic isotope exists. Currently 19 other instable isotopes have been discovered.

What are the health effects of arsenic in water?

Arsenic related illness is usually caused by consumption of contaminated drinking water. In the old days it was applied as a poison, because symptoms of arsenic poisoning resemble cholera symptoms, and therefore the intentional factor was shaded.
Arsenic appears to be essential for some plant and animal species. A possible safe dose for humans was calculated. If arsenic is a dietary mineral, this dose would be 15-25 μg. This amount could be absorbed from food without any trouble. The total amount of arsenic in a human body is about 0.5-15 mg. Many arsenic compounds are absorbed 60-90%, but they are also easily excreted. Humans can develop resistance to certain arsenic concentrations. Shortly after absorption arsenic can be found in liver, spleen, lungs and digestive tract. Most arsenic is excreted, and residues may be found in skin, hair, nails, legs and teeth.
Under conditions of prolonged exposure, many organs may be damaged, skin pigmentation may occur, hair may fall out and nail growth may stop.
Toxicity differs between various arsenic compounds, for example, monomethyl arsenic acid and inorganic arsenide have a higher toxicity level than arsenic choline. Acute toxicity is generally higher for inorganic arsenic compounds than for organic arsenic compounds. Oral intake of more than 100 mg is lethal. The lethal dose of arsenic trioxide is 10-180 mg, and for arsenide this is 70-210 mg. The mechanism of toxicity is binding and blocking sulphur enzymes. Symptoms of acute arsenic poisoning are nausea, vomiting, diarrhoea, cyanosis, cardiac arrhythmia, confusion and hallucinations. Symptoms of chronic arsenic poisoning are less specific. These include depression, numbness, sleeping disorders and headaches.
Arsenic related health effects are usually not acute, but mostly encompass cancer, mainly skin cancer. Arsenic may cause low birth weight and spontaneous abortion.
Arsenic in drinking water is an issue of global importance, therefore the legal limit was decreased to 10 μg/L. This legal limit is not met in countries such as Vietnam and Bangladesh, where millions of people consume drinking water with an arsenic content of over 50 μg/L. This problem results in long-term chronic health effects, such as skin disease, skin cancer, and tumours in lungs, bladder, kidneys and liver.

Which water purification technologies can be applied to remove arsenic from water?

Arsenic removal from water can be carried out in different ways. Options include ion exchange, membrane filtration, and iron and aluminium coagulation. Drinking water mainly contains inorganic arsenic (arsenide or arsenate), therefore determining the total arsenic concentration suffices. Distinguishing between different types of arsenic is irrelevant.
Arsenic removal from soils can be achieved by applying ferns that bioaccumulate large arsenic concentrations.

## What is the molar solubility of #"CaF"_2# in water in terms of its #K_(sp)#?

The molar solubility of an insoluble ionic compound tells you how many moles of said compound you can dissolve in one liter of water.

Insoluble ionic compounds do not dissociate completely in aqueous solution, which implies that an equilibrium is established between the undissolved solid and the dissolved ions.

The solubility product constant, #K_(sp)# , essentially tells you how far to the left this equilibrium lies.

In your case, calcium fluoride, #"CaF"_2# , is considered insoluble in aqueous solution. The small amounts of calcium fluoride that do dissociate will produce calcium cations, #"Ca"^(2+)# , and fluoride anions, #"F"^(-)# , in solution

Notice that every mole of calcium fluoride that dissolves produces #1# mole of calcium cations and #color(red)(2)# moles of fluoride anions.

If you take #s# to be the concentration of calcium fluoride that dissociates, i.e. its molar solubility, you can use an ICE table to find the value of #s#

#color(purple)("I")color(white)(aaaacolor(black)(-)aaaaaaaaaaaacolor(black)(0)aaaaaaaaaacolor(black)(0)#
#color(purple)("C")color(white)(aaaacolor(black)(-)aaaaaaaaaacolor(black)((+s))aaaaaacolor(black)((+color(red)(2)s))#
#color(purple)("E")color(white)(aaaacolor(black)(-)aaaaaaaaaaaacolor(black)(s)aaaaaaaaacolor(black)(color(red)(2)s)#

By definition, the solubility product constant will be

This will be equivalent to

Therefore, the molar solubility of calcium fluoride in terms of its #K_(sp)# will be

## Flow batteries for the grid: going organic

Scientists are making tremendous strides toward creating better batteries—storing more energy at lower cost and lasting longer than ever before. The results touch many aspects of our lives, translating to a more resilient electric grid, longer-lasting laptop batteries, more electric vehicles, and greater use of renewable energy from blowing wind, shining sun, or flowing water.

For grid-scale batteries, identifying the right materials and combining them to create a new recipe for energy storage is a critical step in the world’s ability to harness and store renewable energy. The most widely used grid-scale batteries use lithium-ion technology, but those are difficult to customize moment to moment in ways most useful to the grid, and there are safety concerns. Redox flow batteries are a growing alternative however, most use vanadium, which is expensive, not easily available, and prone to price fluctuations. Those traits pose barriers to widespread grid-scale energy storage.

Alternative materials for flow batteries include organic molecules, which are far more available, more environmentally friendly and less expensive than vanadium. But organics haven’t held up well to the demands of flow-battery technology, usually petering out faster than required. Long-term stability of the molecules is important so they maintain their ability to perform chemical reactions for many years.

“These organic materials are made out of the most common materials available—carbon, hydrogen and oxygen,” said Wei Wang, the PNNL scientist who leads the flow battery team. “They are easily available they don’t need to be mined, as substances like vanadium do. This makes them very attractive for grid-scale energy storage.”

In the Science paper, Wang’s team demonstrated that low-cost organic fluorenone is, surprisingly, not only a viable candidate but also a standout performer when it comes to energy storage.

In laboratory testing that mimicked real-world conditions, the PNNL battery operated continuously for 120 days, ending only when other equipment unrelated to the battery itself wore out. The battery went through 1,111 full cycles of charging and discharging—the equivalent of several years of operation under normal circumstances—and lost less than 3 percent of its energy capacity. Other organic-based flow batteries have operated for a much shorter period.

The flow battery the team created is only about 10 square centimeters, about the size of a large postage stamp, and puts out about 500 milliwatts of power, not even enough to power a cell phone camera. But the tiny structure embodies tremendous promise: Its energy density is more than twice that of the vanadium batteries in use today and its chemical components are inexpensive, long lasting and widely available.

### Oxidation of biogenic and water -soluble compounds in aqueous and organic aerosol droplets by ozone : a kinetic and product analysis approach using laser Raman tweezers

M. D. King, K. C. Thompson, A. D. Ward, C. Pfrang and B. R. Hughes, Faraday Discuss., 2008, 137, 173 DOI: 10.1039/B702199B

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## Aqueous solution

An aqueous solution is a solution in which the solvent is water. It is mostly shown in chemical equations by appending (aq) to the relevant chemical formula. For example, a solution of table salt, or sodium chloride (NaCl), in water would be represented as Na +
(aq) + Cl −
(aq) . The word aqueous (which comes from aqua) means pertaining to, related to, similar to, or dissolved in, water. As water is an excellent solvent and is also naturally abundant, it is a ubiquitous solvent in chemistry. Aqueous solution is water with a pH of 7.0 where the hydrogen ions ( H +
) and hydroxide ions ( OH −
) are in Arrhenius balance ( 10 −7 ).

A non-aqueous solution is a solution in which the solvent is a liquid, but is not water. [1] (See also Solvent and Inorganic nonaqueous solvent.)

Substances that are hydrophobic ('water-fearing') do not dissolve well in water, whereas those that are hydrophilic ('water-friendly') do. An example of a hydrophilic substance is sodium chloride. Acids and bases are aqueous solutions, as part of their Arrhenius definitions.

The ability of a substance to dissolve in water is determined by whether the substance can match or exceed the strong attractive forces that water molecules generate between themselves. If the substance lacks the ability to dissolve in water, the molecules form a precipitate.

Reactions in aqueous solutions are usually metathesis reactions. Metathesis reactions are another term for double-displacement that is, when a cation displaces to form an ionic bond with the other anion. The cation bonded with the latter anion will dissociate and bond with the other anion.

Aqueous solutions that conduct electric current efficiently contain strong electrolytes, while ones that conduct poorly are considered to have weak electrolytes. Those strong electrolytes are substances that are completely ionized in water, whereas the weak electrolytes exhibit only a small degree of ionization in water.

Nonelectrolytes are substances that dissolve in water yet maintain their molecular integrity (do not dissociate into ions). Examples include sugar, urea, glycerol, and methylsulfonylmethane (MSM).

When writing the equations of aqueous reactions, it is essential to determine the precipitate. To determine the precipitate, one must consult a chart of solubility. Soluble compounds are aqueous, while insoluble compounds are the precipitate. There may not always be a precipitate.

When performing calculations regarding the reacting of one or more aqueous solutions, in general one must know the concentration, or molarity, of the aqueous solutions. Solution concentration is given in terms of the form of the solute prior to it dissolving.

Aqueous solutions may contain, especially in the alkaline zone or subjected to radiolysis, hydrated atomic hydrogen and hydrated electrons.

## Ksp : Solubility Product

So let's imgaine an experiment in which I have 100 mL water at 25°C and I add solid barium sulfate (BaSO4) 0.0002 g at a time, stirring between each addition and analysing the solution to determine the concentration of barium ions (Ba 2+ (aq)) and sulfate ions (SO4 2- ).

The results of the experiment might look like this:

total mass BaSO4
[Ba 2+ (aq)] mol L -1 [SO4 2- (aq)] mol L -1 Comments
0 0 0 Concentration of ions in solution is increasing as more BaSO4 is added to the solution.
All added BaSO4 is dissolved, no BaSO4 precipitates out of solution.
BaSO4(s) &rarr Ba 2+ (aq) + SO4 2- (aq)
Solution is unsaturated (that is, more BaSO4 can be dissolved the solution).
0.0002 8.6 × 10 -6 8.6 × 10 -6
0.0004 1.7 × 10 -5 1.7 × 10 -5
0.0006 2.6 × 10 -5 2.6 × 10 -5
0.0008 3.4 × 10 -5 3.4 × 10 -5
0.001 3.9 × 10 -5 3.9 × 10 -5 Concentration of ions in solution is constant.
Equilibrium has been established:
BaSO4(s) Ba 2+ (aq) + SO4 2- (aq)
Solution is saturated (adding more BaSO4 causes "excess" BaSO4 to precipitate)
0.0012 3.9 × 10 -5 3.9 × 10 -5
0.0014 3.9 × 10 -5 3.9 × 10 -5
0.0016 3.9 × 10 -5 3.9 × 10 -5

So dissolving BaSO4 in water should be considered an equilibrium reaction:

Let's use our concentration data above to calculate the value of the ion product at each stage of our experiment:

total mass BaSO4
[Ba 2+ (aq)] mol L -1 [SO4 2- (aq)] mol L -1 [Ba 2+ (aq)][SO4 2- (aq)]
(ion product)
0 0 0 0 Value of ion product is increasing.
Equilibrium has not yet been established.
0.0002 8.6 × 10 -6 8.6 × 10 -6 7.4 × 10 -11
0.0004 1.7 × 10 -5 1.7 × 10 -5 2.9 × 10 -10
0.0006 2.6 × 10 -5 2.6 × 10 -5 6.8 × 10 -10
0.0008 3.4 × 10 -5 3.4 × 10 -5 1.2 × 10 -9
0.001 3.9 × 10 -5 3.9 × 10 -5 1.5 × 10 -9 Value of ion product is constant.
Equilibrium has been established:
BaSO4(s) Ba 2+ (aq) + SO4 2- (aq)
Solution is saturated (adding more BaSO4 causes "excess" BaSO4 to precipitate)
0.0012 3.9 × 10 -5 3.9 × 10 -5 1.5 × 10 -9
0.0014 3.9 × 10 -5 3.9 × 10 -5 1.5 × 10 -9
0.0016 3.9 × 10 -5 3.9 × 10 -5 1.5 × 10 -9

Equilibrium law states that the condition for chemical equilibrium is when the value of the mass-action expression, Q, is equal to the value for the equilibrium constant, Kc.
When an ionic solid dissolves to form ions in solution, Q is the ion product, and the equilibrium constant is called a solubility product and is given the symbol Ksp:

For our barium sulfate experiment:

When the value of the ion product is less than the value for the solubility product (Ksp), the solution is unsaturated, we can continue to add more solute and it will all dissolve, that is, no precipitate forms:

When the value of the ion product is greater than the value for the solubility product (Ksp), the solution is supersaturated with respect to one or both of the ions, the composition of the solution is unstable and the equilibrium position shifts to the left to produce a precipitate, reducing the concentration of ions in solution, until equilibrium is established.

 Values for various solubility products, Ksp, are tabulated on the right. Note the tabulated value of Ksp for barium sulfate at 25°C is 1.5 × 10 -9 .

Note that Ksp is an equilibrium constant so it is temperature dependent, and tables of values are produced for a specific temperature (usually 25°C).

We can use the values of Ksp to:

## Example : Calculating the solubility of an ionic compound (MA).

(based on the StoPGoPS approach to problem solving)

Question: Calculate how much silver bromide in moles will dissolve in 1 L of water at 25°C given Ksp = 5.0 x 10 -13 at 25 o C.

calculate moles of silver bromide dissolved in 1 L of water at 25°C:

n(AgBr(aq)) = ? mol (saturated solution)

AgBr(s) Ag + (aq) + Br - (aq)

Ksp = [Ag + (aq)][Br - (aq)]

From balanced chemical equation: [Ag + (aq)] = [Br - (aq)]

let x = [Ag + (aq)] = [Br - (aq)]

Ksp = x 2

For a saturated solution: x = [Ag + (aq)] = [Br - (aq)] = [AgBr(aq)]

[AgBr(aq)] is moles of AgBr dissolved in 1 L of solution.
Assume volume of AgBr is negligible compared to the volume of water, so 1 L of solution = 1 L of water.
[AgBr(aq)] = n(AgBr(aq)) mol ÷ 1 L
n(AgBr(aq)) = [AgBr(aq)]

x = &radicKsp = &radic5.0 x 10 -13 = 7 × 10 -7 mol L -1

[AgBr(aq)] = 7 × 10 -7 mol L -1

7 × 10 -7 moles of AgBr is dissolved in 1 L of solution.

Use your calculated values for concentration to calculate Ksp (checking your original calculations)

Ksp = x 2 = (7 × 10 -7 ) 2 = 5 × 10 -15

Ksp was given as 5 × 10 -15 so we have confidence in our calculations.

The solubility of AgBr at 25 o C is 7 x 10 -7 mol L -1
7 x 10 -7 moles AgBr will dissolve in 1 L of water at 25 o C.

## Example : Calculating the solubility of an ionic compound (MA2)

(based on the StoPGoPS approach to problem solving)

Question: Calculate how much strontium fluoride in moles per litre will dissolve in 1 L of water given Ksp = 2.5 × 10 -9 at 25 o C.

Calculate solubility of strontium flouride in moles per litre

[SrF2(aq)] = ? mol L -1

Use the calculated values for concentrations to calculate Ksp (checking our calculations):

[Sr 2+ (aq)] = x = 8.5 × 10 -4 mol L -1

[F - (aq)] = 2x = 2 × 8.5 × 10 -4 = 1.7 × 10 -3 mol L -1

Ksp = [Sr 2+ (aq)][F - (aq)] 2
= 8.5 × 10 -4 × (1.7 × 10 -3 ) 2
= 2.5 × 10 -9

Ksp was given as 2.5 × 10 -9 in the question, so we are confident our answer is correct.

The solubility of SrF2 at 25 o C is 8.5 × 10 -4 mol L -1
8.5 × 10 -4 mol L -1 moles SrF2 will dissolve in 1 L of water at 25 o C.

Try the drill questions now!

## Example : Deciding whether a precipitate will form

(based on the StoPGoPS approach to problem solving)

Question: Will a precipitate form if 25.0 mL of 1.4 × 10 -9 mol L -1 NaI and 35.0 mL of 7.9 × 10 -7 mol L -1 AgNO3 are mixed?