Identifying halide ions such as chloride (Cl⁻), bromide (Br⁻), and iodide (I⁻) is fundamental in qualitative chemical analysis. The silver nitrate (AgNO₃) test is a widely used method in the Cambridge IGCSE Chemistry curriculum to detect the presence of these halides. This test leverages the distinct precipitation reactions that halides undergo with silver ions, allowing for their straightforward identification and differentiation.

Key Concepts

Understanding Halides

Halides are monoatomic ions formed when halogen atoms gain an electron. The primary halides examined in qualitative analysis are chloride (Cl⁻), bromide (Br⁻), and iodide (I⁻). These ions vary in their chemical behavior, particularly in their reactions with silver nitrate, which is pivotal for their identification.

Silver Nitrate as a Reagent

Silver nitrate (AgNO₃) is a versatile reagent used in various chemical analyses due to its ability to react with halide ions to form insoluble silver halides. The general reaction can be represented as:

$$ AgNO_3 (aq) + X^- (aq) \rightarrow AgX (s) + NO_3^- (aq) $$

where $ X^- $ represents the halide ions Cl⁻, Br⁻, or I⁻.

Precipitation Reactions with Silver Nitrate

When silver nitrate is added to a solution containing halide ions, distinct precipitates form:

Chloride Ion (Cl⁻): Forms white precipitate silver chloride (AgCl)
Bromide Ion (Br⁻): Forms cream precipitate silver bromide (AgBr)
Iodide Ion (I⁻): Forms yellow precipitate silver iodide (AgI)

These precipitates are insoluble in water but exhibit varying solubility in ammonia and other reagents, which aids in their identification.

Solubility of Silver Halides

The solubility of silver halides in aqueous solutions decreases down the group in the periodic table, meaning AgCl is more soluble than AgBr, which in turn is more soluble than AgI. This trend is crucial for the selective precipitation and identification of halides.

Distinctive Colors of Silver Halides

Each silver halide has a characteristic color that serves as a visual indicator:

Halide Ion	Silver Halide Precipitate	Color
Chloride (Cl⁻)	Silver Chloride (AgCl)	White
Bromide (Br⁻)	Silver Bromide (AgBr)	Cream
Iodide (I⁻)	Silver Iodide (AgI)	Yellow

These color differences are instrumental in distinguishing between the halides during qualitative analysis.

Procedure for Testing Halides with Silver Nitrate

The standard procedure involves adding a few drops of silver nitrate solution to the test solution containing the unknown halide ions. Observing the formation and color of the precipitate allows for the identification of the specific halide present.

Prepare the test solutions containing halide ions.
Add a few drops of AgNO₃ solution to each test solution.
Observe the color and nature of the precipitate formed.
Conduct confirmatory tests if necessary to ensure accurate identification.

Confirmatory Tests for Halides

To ensure the accurate identification of halides, additional tests can be performed:

Ammonia Test: Adding dilute ammonia solution to the precipitate can help distinguish between halides based on their solubility.
Sodium Thiosulfate Test: Treating the precipitate with sodium thiosulfate can convert silver halides back into soluble complexes, aiding in confirmation.

Applications of Silver Nitrate Test

Beyond academic purposes, the silver nitrate test has practical applications in various industries:

Environmental Analysis: Detecting chloride ions in water sources to assess water quality.
Pharmaceuticals: Ensuring the purity of pharmaceutical compounds by identifying halide contaminants.
Forensic Science: Analyzing samples for the presence of halide substances in criminal investigations.

Advanced Concepts

Precipitation Equilibrium and Solubility Product (Ksp)

Understanding the solubility of silver halides involves the concept of solubility equilibrium and the solubility product constant (Ksp). The dissolution of silver halides can be represented as: $$ AgX (s) \leftrightarrow Ag^+ (aq) + X^- (aq) $$

The solubility product expression is given by: $$ K_{sp} = [Ag^+][X^-] $$

The Ksp values decrease from AgCl to AgI, indicating decreasing solubility:

AgCl: $ K_{sp} = 1.8 \times 10^{-10} $
AgBr: $ K_{sp} = 5.0 \times 10^{-13} $
AgI: $ K_{sp} = 8.3 \times 10^{-17} $

These values explain why AgCl is more soluble than AgBr and AgI, influencing the choice of reagents and conditions in qualitative analysis.

Selective Precipitation and Masking Agents

Selective precipitation is essential when multiple halides are present in a solution. By carefully controlling the concentration of AgNO₃ and the reaction conditions, specific halides can be precipitated while others remain in solution. Additionally, masking agents may be used to prevent certain ions from precipitating, allowing for sequential identification of halides.

Interference from Other Ions

In real-world samples, the presence of other ions can interfere with the silver nitrate test. For instance, the presence of sulfide ions (S²⁻) can form silver sulfide (Ag₂S), a black precipitate that may obscure the detection of halides. Therefore, it's crucial to account for and mitigate potential interferences through proper sample preparation and testing protocols.

Quantitative Analysis vs. Qualitative Analysis

While the silver nitrate test is primarily qualitative, determining the exact concentration of halide ions requires quantitative methods such as titration. By measuring the amount of AgNO₃ needed to precipitate the halides, one can calculate their concentrations using stoichiometric principles.

Environmental and Health Implications

Halide ions, particularly chloride, play significant roles in environmental chemistry and human health. Elevated chloride levels in water can lead to corrosion of infrastructure and affect water taste. Iodide is essential for thyroid function, while bromide's impact varies based on its concentration and context. Understanding their presence and concentrations through tests like the silver nitrate assay is vital for maintaining environmental and public health standards.

Historical Development of Halide Identification Techniques

The identification of halides has evolved over centuries, with the silver nitrate test being a cornerstone in qualitative analysis. Early chemists recognized the distinct colors of silver halides, leading to the development of systematic tests for ion identification. Advances in analytical chemistry have since built upon these fundamental principles, enhancing accuracy and enabling the detection of ions in complex mixtures.

Integration with Modern Analytical Techniques

Modern analytical techniques such as spectrophotometry and chromatography have integrated traditional qualitative tests like the silver nitrate assay to enhance their capabilities. These integrations allow for more precise measurements and the ability to handle more complex samples, bridging classical chemistry with contemporary analytical methods.

Comparison Table

Halide Ion	Silver Halide Precipitate	Color of Precipitate	Solubility in Ammonia
Chloride (Cl⁻)	Silver Chloride (AgCl)	White	Soluble
Bromide (Br⁻)	Silver Bromide (AgBr)	Cream	Partially Soluble
Iodide (I⁻)	Silver Iodide (AgI)	Yellow	Insoluble

This table highlights the distinct characteristics of each silver halide, facilitating the differentiation of halide ions during qualitative analysis.

Summary and Key Takeaways

The silver nitrate test is essential for identifying halide ions (Cl⁻, Br⁻, I⁻) through precipitation reactions.
Distinct precipitate colors and solubility properties aid in differentiating between chloride, bromide, and iodide.
Understanding solubility product constants (Ksp) is crucial for interpreting precipitation equilibria.
Advanced concepts include selective precipitation, interference management, and the integration of qualitative tests with modern analytical methods.
Applications of the silver nitrate test extend beyond academia to environmental analysis, pharmaceuticals, and forensic science.

Examiner Tip

Tips

Remember the solubility trend of silver halides with the mnemonic "Clowns Bring Icy Lollies" to recall AgCl > AgBr > AgI in solubility. When performing the silver nitrate test, always add AgNO₃ dropwise to control the reaction and observe precipitate formation accurately. For exam success, practice writing balanced precipitation reactions and familiarize yourself with the characteristic colors of each silver halide to quickly identify halides during tests.

Did You Know

Did you know that silver iodide (AgI) is used in cloud seeding to encourage rainfall? This fascinating application leverages AgI's ability to catalyze ice crystal formation in clouds, promoting precipitation. Additionally, silver halides like AgCl were historically crucial in early photography, as they are sensitive to light and form the basis of photographic film. Another interesting fact is that silver nitrate, beyond its use in halide testing, is employed in medical settings to prevent infant eye infections.

Common Mistakes

Error 1: Confusing the solubility order of silver halides. Students often mix up the solubility trend, forgetting that AgCl is more soluble than AgBr and AgI.
Incorrect: AgI > AgBr > AgCl
Correct: AgCl > AgBr > AgI

Error 2: Misidentifying precipitate colors. For example, mistaking AgBr's cream color for AgI's yellow.
Incorrect: Silver bromide (AgBr) is yellow.
Correct: Silver bromide (AgBr) is cream-colored.

FAQ

Why is silver nitrate used for halide detection?

Silver nitrate selectively reacts with halide ions to form insoluble silver halides, enabling easy identification through precipitate formation.

What color is the precipitate formed with iodide ions?

The precipitate formed with iodide ions (AgI) is yellow.

Can silver nitrate detect all halide ions?

Yes, silver nitrate can detect chloride, bromide, and iodide ions by forming distinct precipitates for each.

What are common interferences in the silver nitrate test?

Anions like sulfide ions can interfere by forming other insoluble silver compounds, leading to false positives.

How can the accuracy of the silver nitrate test be improved?

Using indicator dyes, standardizing reaction conditions, and employing instrumental methods like UV-Vis spectroscopy can enhance accuracy.

1. Acids, Bases, and Salts

1.1 Salt Preparation

1.1.1 Making soluble salts by reaction with excess metal

1.1.2 Making soluble salts by reaction with excess base

1.1.3 Making soluble salts by reaction with excess carbonate

1.1.4 Making soluble salts by titration

1.2 Solubility Rules

1.2.1 Solubility of sodium, potassium, ammonium salts

1.2.2 Solubility of nitrates, chlorides, sulfates, carbonates, and hydroxides

1.3 Hydrated and Anhydrous Salts

1.3.1 Definition of hydrated and anhydrous substances

1.4 Insoluble Salts

1.4.1 Making insoluble salts by precipitation

1.5 Water of Crystallization

1.5.1 Definition and examples (CuSO₄·5H₂O, CoCl₂·6H₂O)

1.6 Properties of Acids

1.6.1 Reaction of acids with metals, bases, and carbonates

1.6.2 Effect of acids on litmus, thymolphthalein, and methyl orange

1.7 Properties of Bases

1.7.1 Definition of bases and alkalis

1.7.2 Reactions of bases with acids and ammonium salts

1.7.3 Effect of alkalis on litmus, thymolphthalein, and methyl orange

1.8 Acids and Alkalis in Aqueous Solution

1.8.1 H⁺ ions in acids, OH⁻ ions in alkalis

1.8.2 Compare acidity and alkalinity using pH and indicators

1.9 Neutralization

1.9.1 Reaction of acids with alkalis to form water

1.10 Strong and Weak Acids

1.10.1 Definition of strong and weak acids

1.10.2 Hydrochloric acid as a strong acid (HCl → H⁺ + Cl⁻)

1.10.3 Ethanoic acid as a weak acid (CH₃COOH ⇌ H⁺ + CH₃COO⁻)

1.11 Oxides

1.11.1 Classify oxides as acidic (e.g., SO₂, CO₂) or basic (e.g., CuO, CaO)

1.11.2 Definition and examples of amphoteric oxides (Al₂O₃, ZnO)

2. Experimental Techniques and Chemical Analysis

2.1 Identification of Ions and Gases

2.1.1 Test for nitrate (NO₃⁻) using aluminum and NaOH

2.1.2 Test for sulfite (SO₃²⁻) using acidified KMnO₄

2.1.3 Cation tests using NaOH and NH₃

2.1.4 Gas tests for NH₃, CO₂, Cl₂, H₂, O₂, SO₂

2.1.5 Flame tests for Li⁺, Na⁺, K⁺, Ca²⁺, Ba²⁺, Cu²⁺

2.1.6 Test for carbonate (CO₃²⁻) using acid

2.1.7 Test for halides (Cl⁻, Br⁻, I⁻) using silver nitrate

2.1.8 Test for sulfate (SO₄²⁻) using barium nitrate

2.2 Experimental Design

2.2.1 Measuring time, temperature, mass, volume

2.2.2 Advantages and disadvantages of methods

2.2.3 Definitions of solvent, solute, solution

2.2.4 Definition of saturated solution

2.2.5 Definition of residue and filtrate

2.3 Acid-Base Titrations

2.3.1 Titration procedure using burette and pipette

2.3.2 Identifying endpoint using indicators

2.4 Chromatography

2.4.1 Paper chromatography for separating mixtures

2.4.2 Identifying substances using chromatograms

2.4.3 Rf value calculation

2.5 Separation and Purification

2.5.1 Filtration, crystallization, distillation

2.5.2 Identifying purity using melting and boiling points

3. Chemical Reactions

3.1 Rate of Reaction

3.1.1 Collision theory: kinetic energy of particles

3.1.2 Collision theory: activation energy

3.1.3 How changes in conditions affect reaction rate

3.1.4 How catalysts lower activation energy

3.1.5 Effect of concentration on reaction rate

3.1.6 Effect of gas pressure on reaction rate

3.1.7 Effect of surface area on reaction rate

3.1.8 Effect of temperature on reaction rate

3.1.9 Effect of catalysts on reaction rate

3.1.10 Definition of catalyst and its role

3.1.11 Methods to measure reaction rate

3.1.12 Interpret graphs of reaction rates

3.1.13 Collision theory: particle concentration

3.1.14 Collision theory: frequency of collisions

3.2 Reversible Reactions and Equilibrium

3.2.1 Definition of reversible reactions

3.2.2 Effect of conditions on reaction direction

3.2.3 Definition of dynamic equilibrium

3.2.4 Effect of temperature on equilibrium

3.2.5 Effect of pressure on equilibrium

3.2.6 Effect of concentration on equilibrium

3.2.7 Effect of catalyst on equilibrium

3.2.8 Haber process equation (N₂ + 3H₂ ⇌ 2NH₃)

3.2.9 Raw materials for Haber process

3.2.10 Conditions used in Haber process

3.2.11 Contact process equation (2SO₂ + O₂ ⇌ 2SO₃)

3.2.12 Raw materials for Contact process

3.2.13 Conditions used in Contact process

3.2.14 Why contact process conditions are used

3.3 Redox

3.3.1 Definition of redox reactions

3.3.2 Oxidation as gain of oxygen, reduction as loss of oxygen

3.3.3 Identify redox reactions using oxygen gain/loss

3.3.4 Definition of oxidation and reduction in terms of electrons

3.3.5 Oxidation and reduction in terms of oxidation numbers

3.3.6 Redox reactions using color changes (MnO₄⁻, I₂)

3.3.7 Definition of oxidizing and reducing agents

3.3.8 Identify oxidizing and reducing agents in reactions

3.3.9 Use of oxidation numbers to identify redox reactions

3.4 Physical and Chemical Changes

3.4.1 Difference between physical and chemical changes

4. Metals

4.1 Corrosion of Metals

4.1.1 Barrier methods (painting, greasing, plastic coating)

4.1.2 How barriers prevent rust (oxygen/water exclusion)

4.1.3 Galvanizing with zinc as protection

4.1.4 Sacrificial protection using reactive metals

4.1.5 Conditions needed for rusting of iron

4.2 Extraction of Metals

4.2.1 Ease of metal extraction depends on reactivity

4.2.2 Blast furnace: iron extracted from hematite

4.2.3 Role of carbon and carbon monoxide in reduction

4.2.4 Thermal decomposition of limestone in furnace

4.2.5 Formation of slag (removes impurities)

4.2.6 Aluminium extraction from bauxite by electrolysis

4.2.7 Blast furnace equations: C + O₂ → CO₂, C + CO₂ → 2CO, Fe₂O₃ + 3CO → 2Fe + 3CO₂

4.2.8 Aluminium extraction: role of cryolite

4.2.9 Why carbon anodes must be replaced in aluminium extraction

4.2.10 Reactions at electrodes in aluminium extraction

4.3 Properties of Metals

4.3.1 Compare physical properties of metals and non-metals

4.3.2 Reactions of metals with acids, water, and oxygen

4.4 Uses of Metals

4.4.1 Uses of aluminium (aircraft, electrical cables, food containers)

4.4.2 Uses of copper (electrical wiring, ductility)

4.5 Alloys

4.5.1 Definition of alloys (mixture of metal with other elements)

4.5.2 Examples: brass (Cu + Zn), stainless steel (Fe + Cr + Ni + C)

4.5.3 Alloys are harder and stronger than pure metals

4.5.4 Why alloys are harder (atoms disrupt layers)

4.6 Reactivity Series

4.6.1 Order of reactivity: K, Na, Ca, Mg, Al, C, Zn, Fe, H, Cu, Ag, Au

4.6.2 Reactions of K, Na, Ca with water

4.6.3 Reaction of Mg with steam

4.6.4 Reactions of Mg, Zn, Fe, Cu, Ag, Au with acids

4.6.5 Explain reactivity based on ion formation

4.6.6 Displacement reactions of metals

4.6.7 Why aluminium appears unreactive (oxide layer)

5. Stoichiometry

5.1 Formulae

5.1.1 Definition of empirical formula

5.1.2 Deduce ionic formula from charges

5.1.3 Write balanced equations with state symbols

5.1.4 Formulae of elements and compounds

5.1.5 Determine formula from atomic ratios

5.1.6 Definition of molecular formula

5.1.7 Write word and symbol equations

5.2 Relative Masses

5.2.1 Definition of relative atomic mass (Ar)

5.2.2 Definition of relative molecular/formula mass (Mr)

5.2.3 Calculate reacting masses (without mole concept)

5.3 The Mole and Avogadro Constant

5.3.1 Concentration in g/dm³ and mol/dm³

5.3.2 Definition of the mole (6.02 × 10²³ particles)

5.3.3 Mole calculations using mass and molar mass

5.3.4 Use molar gas volume (24 dm³ at r.t.p.)

5.3.5 Calculate reacting masses and limiting reactants

5.3.6 Titration calculations (moles, volume, concentration)

5.3.7 Calculate empirical and molecular formulae

5.3.8 Calculate percentage yield and purity

6. Organic Chemistry

6.1 Formulae and Functional Groups

6.1.1 Definition of saturated and unsaturated compounds

6.1.2 Draw and interpret displayed formulae

6.1.3 General formulae for alkanes, alkenes, alcohols, carboxylic acids

6.1.4 Definition of functional groups

6.1.5 Definition of a homologous series

6.2 Naming Organic Compounds

6.2.1 Naming and drawing methane, ethane, ethene, ethanol, ethanoic acid

6.2.2 Identifying organic compounds based on name, formula, or structure

6.3 Fuels

6.3.1 Names of fossil fuels: coal, natural gas, petroleum

6.3.2 Methane as the main component of natural gas

6.3.3 Definition of hydrocarbons

6.3.4 Petroleum as a mixture of hydrocarbons

6.3.5 Fractional distillation of petroleum

6.3.6 Properties of fractions (volatility, boiling points, viscosity)

6.3.7 Uses of petroleum fractions (refinery gas, gasoline, naphtha, kerosene, diesel, fuel oil, lubricatin

6.4 Alkanes

6.4.1 Single covalent bonds in alkanes (saturated hydrocarbons)

6.4.2 Alkanes are generally unreactive except for combustion and substitution

6.5 Alkenes

6.5.1 Double carbon-carbon bonds in alkenes (unsaturated hydrocarbons)

6.5.2 Cracking of large alkanes to produce alkenes and hydrogen

6.5.3 Test for unsaturation using bromine water

6.6 Alcohols

6.6.1 Ethanol production by fermentation of glucose

6.6.2 Ethanol production by catalytic hydration of ethene

6.6.3 Combustion of ethanol

6.6.4 Uses of ethanol (solvent, fuel)

6.7 Carboxylic Acids

6.7.1 Reactions of ethanoic acid with metals, bases, and carbonates

6.7.2 Formation of ethanoic acid by oxidation of ethanol

6.7.3 Ester formation from carboxylic acids and alcohols

6.8 Polymers

6.8.1 Definition of polymers and monomers

6.8.2 Formation of poly(ethene) by addition polymerization

6.8.3 Plastics are made from polymers

6.8.4 Environmental issues of plastics (landfills, ocean waste, toxic gases)

7. States of Matter

7.1 Solids, Liquids, and Gases

7.1.1 Properties of solids, liquids, and gases

7.1.2 Particle arrangement in solids, liquids, and gases

7.1.3 Changes of state: melting, boiling, evaporation, freezing, condensation

7.1.4 Effect of temperature and pressure on gas volume

7.2 Diffusion

7.2.1 Diffusion explained using kinetic particle theory

8. The Periodic Table

8.1 Arrangement of Elements

8.1.1 Periodic Table arranged by atomic number

8.1.2 Change from metallic to non-metallic across a period

8.1.3 Relation of group number to ion charge

8.1.5 Predicting properties using Periodic Table position

8.2 Group I: Alkali Metals

8.2.1 Properties of lithium, sodium, and potassium

8.2.2 Trends down Group I: melting point, density, reactivity

8.2.3 Predict properties of other Group I elements

8.3 Group VII: Halogens

8.3.1 Properties of chlorine, bromine, iodine

8.3.2 Trends down Group VII: density, reactivity

8.3.3 Appearance of halogens at room temperature

8.3.4 Displacement reactions of halogens

8.3.5 Predict properties of other Group VII elements

8.4 Transition Elements

8.4.1 Properties: high density, melting points, colored compounds, catalysts

8.4.2 Variable oxidation states (iron(II) and iron(III))

8.5 Group VIII: Noble Gases

8.5.1 Noble gases are unreactive and monatomic

8.5.2 Explanation of noble gas stability using electron configuration

9. Atoms, Elements, and Compounds

9.1 Elements, Compounds, and Mixtures

9.1.1 Difference between elements, compounds, and mixtures

9.2 Atomic Structure

9.2.1 Structure of an atom (nucleus, electrons, shells)

9.2.2 Relative charges and masses of subatomic particles

9.2.3 Definition of atomic number (proton number)

9.2.4 Definition of mass number (nucleon number)

9.2.5 Electronic configuration for elements (1-20)

9.2.6 Full outer shell in noble gases (Group VIII)

9.2.7 Relation of outer electrons to group number

9.2.8 Relation of electron shells to period number

9.3 Isotopes

9.3.1 Definition of isotopes

9.3.2 Symbols of isotopes (e.g., ₆¹²C, ₁₇³⁵Cl⁻)

9.3.3 Why isotopes have the same chemical properties

9.3.4 Calculating relative atomic mass from isotopes

9.4 Ions and Ionic Bonds

9.4.1 Formation of cations and anions

9.4.2 Definition of ionic bonding

9.4.3 Formation of ionic bonds (dot-and-cross diagrams)

9.4.4 Properties of ionic compounds

9.4.5 Giant lattice structure in ionic compounds

9.5 Simple Molecules and Covalent Bonds

9.5.1 Definition of covalent bonding

9.5.2 Formation of simple covalent molecules (H₂, Cl₂, H₂O, CH₄, NH₃, HCl)

9.5.3 Properties of simple molecular compounds

9.5.4 Dot-and-cross diagrams for CH₃OH, C₂H₄, O₂, CO₂, N₂

9.6 Giant Covalent Structures

9.6.1 Structure of graphite and diamond

9.6.2 Uses of graphite and diamond

9.6.3 Structure of silicon(IV) oxide (SiO₂)

9.7 Metallic Bonding

9.7.1 Definition of metallic bonding

9.7.2 Properties of metals (conductivity, malleability, ductility)

10. Chemistry of the Environment

10.1 Water

10.1.1 Test for water using cobalt(II) chloride and copper(II) sulfate

10.1.2 Test for purity of water using melting and boiling points

10.1.3 Why distilled water is used in chemistry

10.1.4 Substances in natural water (oxygen, metals, plastics, sewage, microbes, fertilizers)

10.1.5 Harmful substances in water (toxins, nitrates, phosphates, microbes)

10.1.6 Water treatment: sedimentation and filtration

10.1.7 Water treatment: activated carbon for odor removal

10.1.8 Water treatment: chlorination to kill microbes

10.2 Fertilizers

10.2.1 Use of ammonium salts and nitrates as fertilizers

10.2.2 NPK fertilizers for nitrogen, phosphorus, potassium supply

10.3 Air Quality and Climate

10.3.1 Composition of clean, dry air (78% N₂, 21% O₂, trace gases)

10.3.2 Sources of CO₂ (fuel combustion), CO (incomplete combustion), CH₄ (decomposition, digestion)

10.3.3 Sources of NOx (car engines), SO₂ (fossil fuels)

10.3.4 Effects of CO₂ and CH₄ (global warming, climate change)

10.3.5 Effects of CO (toxic), particulates (respiratory issues)

10.3.6 Effects of NOx and SO₂ (acid rain, respiratory problems)

10.3.7 How CO₂ and CH₄ cause global warming (heat absorption)

10.3.8 Ways to reduce climate change (trees, hydrogen, renewables)

10.3.9 Reducing acid rain (catalytic converters, low-sulfur fuels, flue gas desulfurization)

10.3.10 Photosynthesis reaction: CO₂ + H₂O → glucose + O₂

10.3.11 Photosynthesis equation: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂

10.3.12 How NOx forms in car engines

10.3.13 How catalytic converters remove NOx (2CO + 2NO → 2CO₂ + N₂)

11. Electrochemistry

11.1 Electrolysis

11.1.1 Definition of electrolysis

11.1.2 Identify anode, cathode, and electrolyte

11.1.3 Electrolysis of molten lead(II) bromide

11.1.4 Electrolysis of concentrated sodium chloride

11.1.5 Electrolysis of dilute sulfuric acid

11.1.6 Products formed at electrodes

11.1.7 Predict products of binary molten compounds

11.1.8 Purpose of electroplating

11.1.9 Process of electroplating metals

11.1.10 Charge transfer in electrolysis (electrons in external circuit)

11.1.11 Charge transfer in electrolysis (ions in electrolyte)

11.1.12 Electrolysis of aqueous copper(II) sulfate (graphite vs copper electrodes)

11.1.13 Predict products for electrolysis of halide solutions

11.1.14 Write ionic half-equations for anode and cathode reactions

11.2 Hydrogen–Oxygen Fuel Cells

11.2.1 Definition of hydrogen–oxygen fuel cell

11.2.2 Advantages and disadvantages of fuel cells vs gasoline engines

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Test for Halides (Cl⁻, Br⁻, I⁻) Using Silver Nitrate

Introduction