Unlocking the secrets of chemical reactions with types of reactions worksheet answers pdf! Dive into the fascinating world of synthesis, decomposition, single and double replacement reactions. Prepare to be amazed by the magic of balancing chemical equations and deciphering the language of atoms. This guide is your key to mastering chemical transformations.
This resource provides a comprehensive overview of chemical reactions, explaining the fundamental concepts, and offering detailed examples to aid in understanding. It walks you through identifying reactants and products, balancing equations, and understanding various reaction types. From simple introductions to complex problem-solving, this guide ensures a solid foundation in chemical reactions.
Introduction to Chemical Reactions
Chemistry is a fascinating world where matter undergoes transformations. These transformations, called chemical reactions, involve the rearrangement of atoms to form new substances with different properties. Understanding these reactions is crucial for comprehending the world around us, from the food we eat to the medicines we take.
Defining Chemical Reactions
Chemical reactions are processes where substances rearrange their constituent atoms to create new substances with unique characteristics. This rearrangement is fundamentally different from physical changes, which alter the form or state of a substance without changing its underlying composition. For instance, melting ice is a physical change; water remains water. Burning wood, however, is a chemical reaction; the wood transforms into ash and gases.
Distinguishing Chemical Reactions from Physical Changes
Key differences between chemical and physical changes lie in the fundamental alteration of the substance’s composition. In physical changes, the molecules stay intact, while in chemical reactions, bonds between atoms are broken and reformed, resulting in entirely new molecules. Indicators of a chemical reaction include changes in color, temperature, formation of a precipitate (solid), production of gas, or light emission.
Types of Chemical Reactions
Chemical reactions manifest in various ways, categorized into distinct types. These categories help us understand the patterns and mechanisms behind these transformations. Common types include synthesis, decomposition, single replacement, and double replacement reactions.
Synthesis Reactions
Synthesis reactions involve combining two or more reactants to form a single product. A simple example is the formation of water from hydrogen and oxygen. This process involves the combination of elements to form a compound.
Decomposition Reactions
Decomposition reactions are the opposite of synthesis reactions. They involve breaking down a single reactant into two or more products. A classic example is the decomposition of water into hydrogen and oxygen through electrolysis.
Single Replacement Reactions, Types of reactions worksheet answers pdf
In single replacement reactions, one element in a compound is replaced by another element. This often occurs when a more reactive element displaces a less reactive element from a compound. For instance, zinc displacing copper in copper sulfate solution.
Double Replacement Reactions
Double replacement reactions involve the exchange of ions between two compounds. This results in the formation of two new compounds, often including the formation of a precipitate or the release of a gas. A common example is the reaction between sodium chloride and silver nitrate, resulting in the formation of silver chloride precipitate.
| Type of Reaction | Description | General Formula | Example |
|---|---|---|---|
| Synthesis | Two or more reactants combine to form a single product. | A + B → AB | 2H2 + O2 → 2H2O |
| Decomposition | A single reactant breaks down into two or more products. | AB → A + B | 2H2O → 2H2 + O2 |
| Single Replacement | One element replaces another element in a compound. | A + BC → AC + B | Zn + CuSO4 → ZnSO4 + Cu |
| Double Replacement | The exchange of ions between two compounds. | AB + CD → AD + CB | NaCl + AgNO3 → NaNO3 + AgCl |
Identifying Reactants and Products
Unveiling the secrets of chemical transformations hinges on understanding the roles of reactants and products. These fundamental components are the actors in every chemical drama, dictating the outcome and driving the changes. Chemical equations are like scripts, meticulously detailing the participants and their transformations.Chemical equations provide a concise and standardized way to represent chemical reactions. They use chemical formulas to represent substances and symbols to indicate the reaction process.
By dissecting these equations, we can pinpoint the reactants, the starting materials, and the products, the resulting substances. This understanding is crucial for predicting and interpreting the course of chemical reactions.
Understanding Reactants and Products
Chemical reactions involve the transformation of reactants into products. Reactants are the substances that undergo change, while products are the substances formed as a result of the reaction. The key to identifying them lies in the arrangement within the equation.
Conventions in Chemical Equations
Chemical equations employ specific conventions to convey information precisely. Coefficients, numerical multipliers placed before chemical formulas, specify the relative amounts of reactants and products. Subscripts, numbers written below symbols, indicate the number of atoms of each element within a molecule. Understanding these conventions is paramount for accurate interpretation.
Examples of Balanced Chemical Equations
Here are a few examples of balanced chemical equations, highlighting the different components.
2H2 + O 2 → 2H 2O
In this equation, hydrogen (H 2) and oxygen (O 2) are the reactants, and water (H 2O) is the product. The coefficients indicate the mole ratio in which these substances react.
CH4 + 2O 2 → CO 2 + 2H 2O
Methane (CH 4) reacts with oxygen (O 2) to yield carbon dioxide (CO 2) and water (H 2O). The coefficients show the balanced ratio of reactants and products.
Reaction Types and Their Reactants/Products
Different types of reactions exhibit distinct patterns in their reactant and product relationships. This table illustrates examples of common reaction types.
| Reaction Type | Reactants | Products |
|---|---|---|
| Synthesis | 2Na + Cl2 | 2NaCl |
| Decomposition | 2H2O2 | 2H2O + O2 |
| Single Replacement | Zn + CuSO4 | ZnSO4 + Cu |
| Double Replacement | AgNO3 + NaCl | AgCl + NaNO3 |
| Combustion | C3H8 + 5O2 | 3CO2 + 4H2O |
This table showcases the variety of reactions and their characteristic reactants and products. Each reaction type has a unique pattern that allows us to anticipate the products.
Types of Chemical Reactions
Chemistry is the study of matter and its transformations. Chemical reactions are the heart of this transformation, and understanding the different types helps us predict and control these processes. From the explosive combustion of fireworks to the slow rusting of metal, these reactions are everywhere. Knowing the types of reactions is like having a roadmap to navigate the chemical world.
Synthesis Reactions (Combination Reactions)
Synthesis reactions, also known as combination reactions, involve two or more substances combining to form a single product. This is a fundamental type of reaction where simpler substances merge into more complex ones. Think of it like building blocks – combining individual blocks to create a larger, more intricate structure.
- Synthesis reactions are characterized by the formation of a more complex molecule from simpler ones.
- These reactions are represented by a general formula: A + B → AB, where A and B are the reactants and AB is the product.
- Examples include the reaction of hydrogen gas with oxygen gas to produce water: 2H 2(g) + O 2(g) → 2H 2O(l). Another example is the reaction of magnesium metal with oxygen gas to produce magnesium oxide: 2Mg(s) + O 2(g) → 2MgO(s).
Predicting Products of Synthesis Reactions
To predict the products of synthesis reactions, consider the elements involved. Knowing the common oxidation states of elements helps in determining the formula of the product. For instance, if you combine a metal and a nonmetal, the resulting product will likely be an ionic compound. The metal will lose electrons to form a positive ion, and the nonmetal will gain electrons to form a negative ion.
These oppositely charged ions will then attract each other to form the ionic compound. Consider the reaction of sodium (Na) with chlorine (Cl 2). Sodium readily loses one electron to form Na +, while chlorine readily gains one electron to form Cl –. The resulting ionic compound is sodium chloride (NaCl).
Decomposition Reactions
Decomposition reactions are the reverse of synthesis reactions. Instead of combining substances to form a more complex product, decomposition reactions break down a complex substance into simpler substances. This is like taking apart a larger structure into its individual building blocks.
- Decomposition reactions involve a single reactant breaking down into two or more products.
- These reactions are represented by a general formula: AB → A + B, where AB is the reactant and A and B are the products.
- Examples include the decomposition of water into hydrogen and oxygen: 2H 2O(l) → 2H 2(g) + O 2(g). Another example is the decomposition of calcium carbonate into calcium oxide and carbon dioxide: CaCO 3(s) → CaO(s) + CO 2(g).
Predicting Products of Decomposition Reactions
Predicting products of decomposition reactions depends on the nature of the reactant. Consider the type of bonds within the reactant. Ionic compounds, for example, often decompose into their constituent ions. Covalent compounds might decompose into simpler molecules. The decomposition of a compound depends on the conditions under which it is heated or decomposed.
Single Replacement Reactions (Single Displacement Reactions)
Single replacement reactions involve an element replacing another element in a compound. This is a bit like a substitution, where one element takes the place of another in a chemical structure. The ability of an element to displace another depends on the reactivity of the elements involved. A key concept here is the activity series. The activity series lists elements in order of their decreasing reactivity.
A more reactive element can displace a less reactive element from a compound.
| Type of Reaction | General Formula | Examples | Products |
|---|---|---|---|
| Synthesis | A + B → AB | 2Mg + O2 → 2MgO | Magnesium Oxide |
| Decomposition | AB → A + B | 2H2O → 2H2 + O2 | Hydrogen and Oxygen |
| Single Replacement | A + BC → AC + B | Cu + 2AgNO3 → Cu(NO3)2 + 2Ag | Copper(II) Nitrate and Silver |
Double Replacement Reactions
Double replacement reactions, also known as metathesis reactions, are a common type of chemical reaction where the positive and negative ions of two ionic compounds exchange partners. Imagine two dance partners swapping partners, leading to a new pairing. This swapping results in the formation of new compounds. These reactions often occur in aqueous solutions and are driven by the formation of a solid precipitate, the release of a gas, or the creation of a weaker electrolyte.
Understanding the Process
Double replacement reactions typically involve two ionic compounds dissolved in a solution. These compounds dissociate into their constituent ions. The positive ions from one compound swap places with the positive ions from the other compound. This exchange leads to the formation of two new compounds. One of these new compounds might be a solid precipitate, a gas, or a weak electrolyte.
The driving force behind this process is the formation of a product that is less soluble or more stable than the reactants.
Examples of Double Replacement Reactions
Here are some illustrative examples of double replacement reactions:
- Silver nitrate (AgNO 3) reacts with sodium chloride (NaCl) to form silver chloride (AgCl), a white precipitate, and sodium nitrate (NaNO 3). This is a classic example of a precipitation reaction.
- Hydrochloric acid (HCl) reacts with sodium hydroxide (NaOH) to form sodium chloride (NaCl) and water (H 2O). This reaction is an example of acid-base neutralization, resulting in a neutral solution.
- Calcium carbonate (CaCO 3) reacts with hydrochloric acid (HCl) to form calcium chloride (CaCl 2), carbon dioxide gas (CO 2), and water (H 2O). The evolution of carbon dioxide gas is a clear indication of this type of reaction.
Conditions for Double Replacement Reactions
Double replacement reactions often occur under specific conditions, leading to different observable outcomes.
- Precipitation: One of the products is an insoluble solid, called a precipitate. The precipitate forms due to the low solubility of the product in the solvent. This is often indicated by a visible solid separating from the solution.
- Gas Formation: One of the products is a gas. The gas often escapes from the solution, which can be observed as bubbles. The gas formation is often due to the instability of the product in the solution.
- Acid-Base Neutralization: One of the reactants is an acid, and the other is a base. The products of this reaction are a salt and water. The reaction results in a decrease in acidity and basicity of the solution.
Summary Table
| Reaction | Reactants | Products | Observations |
|---|---|---|---|
| Precipitation | AgNO3(aq) + NaCl(aq) | AgCl(s) + NaNO3(aq) | White precipitate forms |
| Gas Formation | CaCO3(s) + 2HCl(aq) | CaCl2(aq) + CO2(g) + H2O(l) | Bubbles of carbon dioxide gas evolve |
| Acid-Base Neutralization | HCl(aq) + NaOH(aq) | NaCl(aq) + H2O(l) | No visible change, but the solution becomes neutral |
Balancing Chemical Equations: Types Of Reactions Worksheet Answers Pdf
Chemical reactions are like recipes for transforming substances. Just like a recipe needs to have the right amounts of ingredients to make a delicious dish, chemical reactions need a precise balance of atoms to proceed correctly. Understanding and accurately representing these balances is crucial in chemistry.Balancing chemical equations is essential for several reasons. It reflects the fundamental law of conservation of mass, a cornerstone of chemistry.
An unbalanced equation incorrectly suggests that atoms are created or destroyed during a reaction, which is impossible. Furthermore, it allows for precise calculations in stoichiometry, enabling chemists to determine the amounts of reactants and products involved.
Importance of Balancing Chemical Equations
Balancing chemical equations is critical for understanding and predicting the outcomes of chemical transformations. An accurate representation ensures that the number of atoms of each element is the same on both sides of the equation, adhering to the law of conservation of mass. This accuracy is fundamental for precise calculations in stoichiometry, allowing for reliable predictions about the amounts of reactants and products involved in a reaction.
Method of Balancing Chemical Equations by Inspection
The method of inspection involves systematically adjusting coefficients in front of chemical formulas until the number of atoms of each element is equal on both sides of the equation. This approach, while seemingly trial-and-error, often becomes quite intuitive with practice. Start by focusing on the elements that appear in more than one compound on either side of the equation.
Examples of Balanced Chemical Equations
- For a combustion reaction of methane (CH 4):
CH4(g) + 2O 2(g) → CO 2(g) + 2H 2O(g)
In this example, the balanced equation shows that one molecule of methane reacts with two molecules of oxygen to produce one molecule of carbon dioxide and two molecules of water.
- For a single displacement reaction of zinc with hydrochloric acid:
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H 2(g)
Here, one atom of zinc reacts with two molecules of hydrochloric acid to produce one molecule of zinc chloride and one molecule of hydrogen gas.
- For a double replacement reaction of silver nitrate and sodium chloride:
AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO 3(aq)
In this reaction, one molecule of silver nitrate reacts with one molecule of sodium chloride to produce one molecule of silver chloride and one molecule of sodium nitrate.
Comparing Balanced and Unbalanced Equations
| Feature | Balanced Equation | Unbalanced Equation |
|---|---|---|
| Atoms of each element | Equal on both sides of the equation | Unequal on both sides of the equation |
| Mass conservation | Adheres to the law of conservation of mass | Violates the law of conservation of mass |
| Stoichiometry | Allows for accurate stoichiometric calculations | Leads to inaccurate stoichiometric calculations |
| Predictive power | Provides accurate predictions about the reaction | Provides inaccurate predictions about the reaction |
Worksheet Structure and Examples
Unveiling the secrets of chemical reactions is like cracking a code! These worksheets are your keys to understanding the different ways atoms rearrange themselves. They’ll guide you through the world of reactants, products, and the essential balancing act of equations.Understanding the structure of a reaction worksheet is crucial for effective learning. A well-structured worksheet will help you systematically analyze chemical reactions, making the process easier and more engaging.
These worksheets often follow a consistent pattern, allowing you to master the concepts more efficiently.
Worksheet Structure
A typical worksheet on reaction types generally starts with clear definitions and explanations of the different reaction types, like synthesis, decomposition, single replacement, and double replacement. This initial section serves as a solid foundation. Next, you’ll typically encounter examples of each type of reaction, along with the necessary steps to balance the chemical equations. This structured approach ensures a comprehensive understanding.
Question Formats
Questions related to reactants and products usually involve identifying the substances involved in a given reaction. For instance, a question might ask to identify the reactants and products in a specific chemical equation. Another common question type deals with balancing chemical equations, requiring students to determine the correct coefficients to ensure that the number of atoms of each element is the same on both sides of the equation.
These types of questions are designed to build a strong grasp of reaction mechanisms.
Sample Worksheet Question
Identify the reactants and products in the following reaction and balance the equation:
H2 + O 2 → H 2O
Answer
Reactants: H 2 (hydrogen) and O 2 (oxygen); Products: H 2O (water).Balanced Equation: 2H 2 + O 2 → 2H 2O
Worksheet Problem Examples
| Reaction | Reactants | Products | Balanced Equation |
|---|---|---|---|
| Synthesis of Water | Hydrogen (H2) and Oxygen (O2) | Water (H2O) | 2H2 + O2 → 2H2O |
| Decomposition of Calcium Carbonate | Calcium Carbonate (CaCO3) | Calcium Oxide (CaO) and Carbon Dioxide (CO2) | CaCO3 → CaO + CO2 |
| Single Replacement of Zinc and Copper Sulfate | Zinc (Zn) and Copper Sulfate (CuSO4) | Zinc Sulfate (ZnSO4) and Copper (Cu) | Zn + CuSO4 → ZnSO4 + Cu |
These examples demonstrate the types of problems often found on reaction worksheets. Mastering these concepts is essential for navigating the fascinating world of chemistry.
Common Mistakes and Troubleshooting
Navigating the world of chemical reactions can sometimes feel like trying to assemble a complicated Lego set without the instructions. Understanding common pitfalls and how to fix them is key to mastering this fascinating field. This section will equip you with the tools to identify and overcome these challenges, ensuring you confidently classify reaction types.
Identifying Common Errors
Students often encounter difficulties in distinguishing between different reaction types. Misinterpretations of reactant and product relationships are a frequent source of errors. For example, failing to recognize the presence of a specific reactant can lead to misclassifying a reaction. Similarly, a misunderstanding of the overall energy changes in a reaction can obscure its type. Careful observation of the reactants and products, along with an understanding of the expected energy changes, are crucial for accurate identification.
Misconceptions About Reaction Types
A common misconception is that all reactions involve a single, obvious change. In reality, some reactions may exhibit subtle transformations that aren’t immediately apparent. Another misconception involves overlooking the role of catalysts. Catalysts, often overlooked, can significantly influence the rate of a reaction without being consumed themselves. Recognizing the subtle signs of a reaction, including the presence of catalysts, is critical for correct classification.
Troubleshooting Steps
Troubleshooting reaction type identification involves a systematic approach. First, thoroughly examine the reactants and products. Second, identify the key characteristics of the reaction. Third, compare the observed characteristics to the definitions of the various reaction types. If the reaction doesn’t fit any known type, consider the possibility of an unusual or complex reaction mechanism.
Troubleshooting Table
| Common Mistake | Explanation | Effective Solution |
|---|---|---|
| Ignoring the presence of catalysts | Failing to acknowledge the influence of catalysts on reaction rates. | Carefully examine the reaction conditions and identify any catalysts present. Review the role of catalysts in the reaction process. |
| Misinterpreting the energy changes | Ignoring the energy changes in the reaction. | Consider the overall energy changes during the reaction. Understand how the energy changes correlate with the different reaction types. |
| Inability to identify the specific reactants | Difficulty in accurately identifying all reactants in a reaction. | Carefully analyze the chemical equation, identifying all substances involved. Review the balanced equation. |
| Overlooking the role of intermediates | Not recognizing the role of intermediates in complex reactions. | Understand that some reactions may involve multiple steps with intermediate products. Analyze the overall reaction equation. |
Illustrative Examples
Let’s dive into the fascinating world of chemical reactions! Understanding these transformations is key to comprehending the universe around us, from the tiniest particles to the grandest explosions. We’ll explore real-world examples of synthesis, decomposition, single replacement, and double replacement reactions, highlighting the crucial role of stoichiometry.
Synthesis Reactions
Synthesis reactions, also known as combination reactions, involve two or more substances combining to form a single product. These reactions often release energy in the form of heat or light. A classic example is the formation of water from hydrogen and oxygen.
2H2(g) + O 2(g) → 2H 2O(l)
In this reaction, two molecules of hydrogen gas combine with one molecule of oxygen gas to produce two molecules of liquid water. Notice how the number of atoms of each element is balanced on both sides of the equation. This balance is crucial to understanding the quantitative relationships in the reaction.
Decomposition Reactions
Decomposition reactions are the reverse of synthesis reactions. A single reactant breaks down into two or more simpler substances. A common example is the decomposition of calcium carbonate (limestone).
CaCO3(s) → CaO(s) + CO 2(g)
Heating calcium carbonate leads to the formation of calcium oxide (lime) and carbon dioxide gas. This reaction is crucial in various industrial processes, like the production of lime for construction.
Single Replacement Reactions, Types of reactions worksheet answers pdf
Single replacement reactions involve an element replacing another element in a compound. A common example is the reaction of zinc metal with hydrochloric acid.
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H 2(g)
In this reaction, zinc metal displaces hydrogen from hydrochloric acid, forming zinc chloride solution and hydrogen gas. This reaction demonstrates the reactivity series, where zinc is more reactive than hydrogen.
Double Replacement Reactions
Double replacement reactions, also known as metathesis reactions, involve the exchange of ions between two compounds. A common example is the reaction of sodium chloride (table salt) with silver nitrate.
NaCl(aq) + AgNO3(aq) → NaNO 3(aq) + AgCl(s)
This reaction results in the formation of a precipitate of silver chloride, a white solid, and sodium nitrate solution. Predicting the products and states of matter is essential in understanding double replacement reactions.
Stoichiometry in Reaction Equations
Stoichiometry is the quantitative study of reactants and products in chemical reactions. The coefficients in a balanced chemical equation represent the molar ratios of reactants and products. For instance, in the synthesis of water, the coefficients show that 2 moles of hydrogen react with 1 mole of oxygen to produce 2 moles of water. This ratio is essential for calculating the amounts of reactants needed or products formed in a reaction.
Table of Reaction Types
| Reaction Type | Illustration/Diagram | Description |
|---|---|---|
| Synthesis | Two or more reactants combine to form a single product. | (Imagine two different colored balls joining to form a larger ball of a new color) |
| Decomposition | A single reactant breaks down into two or more products. | (Imagine a large ball splitting into two or more smaller balls of different colors) |
| Single Replacement | An element replaces another element in a compound. | (Imagine a smaller ball pushing a larger ball out of a larger ball) |
| Double Replacement | Two compounds exchange ions to form two new compounds. | (Imagine two larger balls exchanging smaller balls with each other) |
Practice Problems
Embark on a thrilling chemical reaction adventure! These practice problems will test your understanding of different reaction types, from simple combinations to intricate double replacements. Get ready to apply your knowledge and unlock the secrets of chemical transformations.
These problems are designed to progressively increase in complexity, ensuring you master each concept before moving on. The solutions are provided to aid your understanding and celebrate your progress. This will be a fantastic opportunity to refine your skills and solidify your grasp on chemical reactions.
Balancing Chemical Equations
Chemical equations, like carefully crafted recipes, must balance the number of atoms on both sides of the equation. This ensures that the law of conservation of mass is upheld, meaning atoms are neither created nor destroyed in a chemical reaction. The following examples showcase the importance of balancing equations.
- Problem 1: Balance the equation: C 3H 8(g) + O 2(g) → CO 2(g) + H 2O(g)
- Solution 1: C 3H 8(g) + 5O 2(g) → 3CO 2(g) + 4H 2O(g)
- Problem 2: Balance the equation: Fe 2O 3(s) + CO(g) → Fe(s) + CO 2(g)
- Solution 2: Fe 2O 3(s) + 3CO(g) → 2Fe(s) + 3CO 2(g)
Single Replacement Reactions, Types of reactions worksheet answers pdf
Single replacement reactions, often referred to as single displacement reactions, involve one element replacing another in a compound. These reactions are common in various applications, from industrial processes to everyday chemistry.
- Problem 3: Predict the products of the following reaction: Zn(s) + CuSO 4(aq) → ?
- Solution 3: Zn(s) + CuSO 4(aq) → ZnSO 4(aq) + Cu(s)
- Problem 4: Predict the products and balance the equation: Al(s) + HCl(aq) → ?
- Solution 4: 2Al(s) + 6HCl(aq) → 2AlCl 3(aq) + 3H 2(g)
Double Replacement Reactions
Double replacement reactions, also known as double displacement reactions, involve the exchange of ions between two compounds. This type of reaction often results in the formation of a precipitate, a gas, or a molecular compound.
- Problem 5: Predict the products and balance the equation: AgNO 3(aq) + NaCl(aq) → ?
- Solution 5: AgNO 3(aq) + NaCl(aq) → AgCl(s) + NaNO 3(aq)
- Problem 6: Predict the products and balance the equation: KOH(aq) + H 2SO 4(aq) → ?
- Solution 6: 2KOH(aq) + H 2SO 4(aq) → K 2SO 4(aq) + 2H 2O(l)
Practice Problems Table
| Problem | Reactants | Products | Balanced Equation |
|---|---|---|---|
| Problem 1 | C3H8(g), O2(g) | CO2(g), H2O(g) | C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g) |
| Problem 2 | Fe2O3(s), CO(g) | Fe(s), CO2(g) | Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g) |
| Problem 3 | Zn(s), CuSO4(aq) | ZnSO4(aq), Cu(s) | Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s) |
| Problem 4 | Al(s), HCl(aq) | AlCl3(aq), H2(g) | 2Al(s) + 6HCl(aq) → 2AlCl3(aq) + 3H2(g) |
| Problem 5 | AgNO3(aq), NaCl(aq) | AgCl(s), NaNO3(aq) | AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq) |
| Problem 6 | KOH(aq), H2SO4(aq) | K2SO4(aq), H2O(l) | 2KOH(aq) + H2SO4(aq) → K2SO4(aq) + 2H2O(l) |
