AP Chemistry 4.1 Introduction for Reactions Study Notes - New Syllabus Effective fall 2024
AP Chemistry 4.1 Introduction for Reactions Study Notes- New syllabus
AP Chemistry 4.1 Introduction for Reactions Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Identify evidence of chemical and physical changes in matter.
Key Concepts:
- Physical & Chemical Changes
- Representing Chemical Changes
4.1.A.1 Physical Change:
1. Defination:
A physical change is a change in the form or appearance of matter, but not in its chemical composition. In other words, the substance remains the same, but its physical properties—like shape, size, state, or texture—are changed.
Some examples of physical changes are:
Melting ice (from solid to liquid)
Boiling water (from liquid to gas)
Crushing a can (change in shape)
Dissolving sugar in water, no new substance has been made
Physical changes are usually reversible, that is, one often can change the substance back into its original form (example: freezing of ice back to water).
2. Phase Changes:
Phase changes (or state changes) refer to the transition of matter from one state to another: solid to liquid, liquid to gas, or gas to solid. They occur by adding or removing heat energy from a substance. Important to know: phase changes are physical changes since only the state changes, not the chemical makeup of the substance.
Below are the key phase changes and the changes between the three states: solid, liquid, and gas:
i. Solid → Liquid (Melting)
Melting is the process in which a solid absorbs heat and changes into a liquid. The molecules in the solid start moving faster, breaking free from their fixed positions.
Example: Ice melting into water.
ii. Liquid → Solid (Freezing or Solidification)
Freezing is the change of heat loss that occurs when a liquid turns into a solid. The molecules of the liquid are slowing down and the distances between the molecules become closer as it becomes a solid.
Example: Water freezing to be ice.
iii. Liquid → Gas (Vaporization)
Vaporization is when a liquid takes in heat and becomes a gas. There are two types of vaporization:
Evaporation: It occurs at the surface of a liquid at any temperature (typically at lower temperatures).
Boiling: It occurs throughout the liquid when it reaches its boiling point.
Example: Water boiling into steam.
iv. Gas → Liquid (Condensation)
Condensation is the process in which a gas loses heat and becomes a liquid. The gas molecules slow down and come together to form a liquid.
Illustration: Condensation of water vapor to liquid water on a chilled glass
v. Solid → Gas (Sublimation)
Sublimation is the change of a solid directly to a gas without passing through the intermediate stage of liquid. This occurs when the solid achieves enough energy for its molecules to transition directly into the gas phase.
Example: Sublimation of dry ice (solid carbon dioxide) to carbon dioxide gas.
vi. Gas → Solid (Deposition)
Deposition is the opposite of sublimation, in which a gas turns directly into a solid without passing through the liquid phase.
Example: Frost forming on a cold surface (water vapor turning directly into ice).
These transitions are dependent on both temperature and pressure. For example, reducing pressure can make it boil at lower temperatures, whereas raising the temperature can make a substance melt or vaporize.
3. Formation and Separation of Mixtures:
The formation and separation of mixtures involve combining or separating substances without altering their individual chemical compositions. Mixtures can be heterogeneous: that is, the components are not uniformly distributed, or homogeneous, where the components are evenly distributed. The formation and separation of mixtures usually depend on the differences in the physical properties of the substances.
Formation of Mixtures:
When two or more substances are combined, they form a mixture, but their chemical identity is not changed. The original properties of the components are preserved and can frequently be separated into their individual parts.
Types of Mixtures:
i. Homogeneous Mixtures (Solutions):
In homogeneous mixtures, the substances are uniformly distributed, and the mixture looks the same throughout.
Illustration: Saltwater is a homogeneous mixture because salt is uniformly dissolved in water.
Process of Formation: This commonly takes place when one substance dissolves in another, such as how sugar dissolves within water.
ii. Heterogeneous Mixtures:
In heterogeneous mixtures, the elements are not mixed and separate, and different parts can be identified or physically isolated.
Example: A salad (leafy greens, tomatoes, and cucumbers) is a heterogeneous mixture because the components retain their individual characteristics and can be picked apart.
Formation Process: This occurs when substances are mixed together but not dissolved or evenly distributed, such as when sand is mixed with gravel.
Separation of Mixtures:
Mixtures can be separated based on their physical properties into the original individual components. These include size, density, solubility, and even magnetic properties. The difference between them is used to achieve the separation process.
Common Methods of Separation:
i. Filtration:
Separate solid particles from liquids or gases through a filter.
Example: Passing a mixture that contains water and sand through a filter so as to separate them.
– How it works:** A filter lets through the small particles, liquid or gas, but captures larger solid particles.
ii. Evaporation:
Removes a solvent from the solution by heating the mixture that makes the solvent evaporate, leaving its solute behind.
Example: Remove the water from saltwater by evaporation, and the salt remains.
How it works: Heat applied to the mixture makes the liquid solvent evaporate and leaves out the solid solute.
iii. Distillation:
Utilized in order to separate a mixture of liquids using their difference in boiling point.
Example: Separating alcohol from the mixture of alcohol and water.
How it works: The mixture is heated, and first, the element that has a lower boiling point vaporizes. Then its vapor is collected after cooling down to condense.
vi. Magnetic Separation:
It is used to separate the magnetic from the non-magnetic.
Example: Separation of iron filings from the mixture of sand and iron using a magnet.
How it works: It attracts magnetic materials and leaves non-magnetic ones behind.
v. Chromatography:
Separation of different components of a mixture, often applied to liquids or gases.
Example: Separation of different pigments in ink.
Working: It flows through a medium (paper or column) and the components travel at different rates, separating according to their affinity for the medium.
vi. Centrifugation:
An application wherein the separation of the components of the mixture is subjected to density, which is spun at high speed.
Example: Separation of blood plasma from blood cells.
How it works: The denser particles go to the bottom of the vessel through centrifugal force.
Key Points to Remember:
In mixtures, the separate substances do not lose their chemical properties and can be separated using physical means.
Homogeneous mixtures (solutions) are uniform throughout; heterogeneous mixtures contain distinguishable, recognizable components.
Separation methods depend on the physical properties, such as size, boiling point, solubility, or magnetic properties.
4.1.A.2 Chemical Change:
1.Definition & Signs: It is a chemical reaction. This is a process in which one or more substances undergo transformations into new substances whose chemical composition and properties are different. In contrast to physical changes, the chemical nature of the original substances is changed in that new molecules or compounds are formed. Chemical changes often are irreversible; that is, it may not be possible to easily recover the original substances by simple physical means.
i. Signs of Chemical Change
There are many indications that a chemical change has occurred. These signs point to the production of a new substance. The most common indications include:
a. Heat: A chemical reaction can release or absorb energy in the form of heat. When it releases heat, it is referred to as an exothermic reaction. When heat is absorbed during the reaction, it is known as an endothermic reaction.
b. Light: Sometimes a chemical reaction can release energy as light, which is noticeable as glowing or brightening.
Example: The reaction between hydrogen and oxygen to form water releases heat and light (combustion).
c. Gas Formation: The formation of a gas is a common indicator of a chemical change. Bubbles or fumes may appear in a liquid as a result of a reaction.
Example: Carbon dioxide gas bubbles form when baking soda reacts with vinegar.
d. Change in Color:
In the formation of a new substance with new properties, it can manifest itself in color changes, particularly with transition between two or more states.
Example: Iron, when it rusts, turns from a metallic gray to reddish-brown color. Similarly, apple is brown when cut and exposed to air as well due to oxidation.
e. Formation of Precipitate:A precipitate is a solid that forms and settles in a liquid during a chemical reaction. This occurs whenever two solutions are mixed and an insoluble substance develops.
Example: A yellow precipitate of lead(II) iodide forms when lead(II) nitrate solution is mixed with potassium iodide solution.
f. Odor Change:Chemical changes may form a new substance that has a particular or different smell.
Example: Rotten eggs smell, often due to hydrogen sulfide gas, results when the right chemicals react with sulfur compounds.
ii. Examples of Chemical Changes:
a. Burning of Wood: The products of the reaction of wood burning include carbon dioxide, water vapor, and ash. This is a chemical change since new substances are formed along with the release of energy in the form of heat and light.
b. Rusting of Iron: Iron reacts with oxygen and water to produce iron oxide, popularly known as rust. This is another chemical change since the original metal has been chemically transformed into a different substance called rust, and the reverse is not possible.
c. Boiling an Egg: When you boil an egg, the proteins in the white portion and yolk change chemically. They become denatured and coagulated in the form of a more solid product. An egg which has been cooked cannot be reversed back to how it was when it was boiled.
d. Digesting Food: Food, during digestion, changes chemically through the work of enzymes, which break down complex molecules into simpler ones. A number of new substances are produced during this period, such as glucose from carbohydrates.
e. Tarnishing of Silver: Silver reacts with the sulfur compounds present in the atmosphere to form a tarnish in the form of silver sulfide. This is a chemical change because the appearance of the silver changes into something new.
iii. Key Differences Between Physical and Chemical Changes:
Physical Change | Chemical Change |
---|---|
No new substances are formed | New substances with different properties are formed |
The chemical composition stays the same | The chemical composition changes |
Generally reversible (e.g., melting, freezing) | Usually irreversible without another reaction |
Changes in form, size, or state of matter | Changes in molecular structure and properties |
2. Energy & Bonding:
Energy change is the fundamental characteristic of chemical reactions. In most cases, it appears in the form of heat, light, or sound. The changes are important factors in determining the way the reaction occurs and if it will be spontaneous or not.
i. Exothermic Reactions:
In exothermic reactions, the energy released when the new bonds are formed is more than that required to break the bonds in the reactants. This results in a net release of energy, usually in the form of heat or light.
Characteristics:
a. Energy is released into the surroundings.
b. The temperature of the surroundings usually increases.
c. These reactions are usually spontaneous.
d. Examples of Exothermic Reactions:
Combustion: Combustion of any fuel like wood or gasoline is an exothermic reaction. It comes with energy release in the form of heat and light.
Respiration: The breakdown of glucose by the body for its energy requirement also represents an exothermic reaction.
Water Formation: Water is produced when hydrogen and oxygen undergo reaction with a huge release of energy.
Example Reaction:
ii. Endothermic Reactions:
For an endothermic reaction, energy is used for breaking the bond within the reactants than being evolved when making the new bond. This process would result in an overall energy absorption, and the surroundings commonly become cooler, too.
Characteristics:
a. This reaction absorbs the energy from the environment.
b. Environment’s temperature becomes lower.
Rearranging, these are not spontaneous reaction types. On the contrary, these require some amount of energy to be given to it.
Examples of Endothermic Reactions
a. Photosynthesis: Plants absorb sunlight (energy) to convert carbon dioxide and water into glucose and oxygen.
b. Dissolving Salt in Water: The dissolution of some salts, such as ammonium nitrate, is endothermic and cools the solution.
c. Thermal Decomposition: Calcium carbonate decomposes into calcium oxide and carbon dioxide by adding heat.
Example Reaction:
(The breakdown of calcium carbonate needs to absorb heat to proceed.)
The decomposition of calcium carbonate requires absorption of heat to take place.
iii. Activation Energy
Every chemical reaction requires a certain amount of energy to start the process, which is called activation energy. It is the energy needed to break the bonds of the reactants so that new bonds can be formed. Even exothermic reactions need this “push” to get started.
Example: An example of activation energy is lighting a match. You have to strike the match so that enough heat is produced to start the chemical reaction (combustion).
iv. Energy Diagrams for Reactions:
An energy diagram, or reaction coordinate diagram, is used to represent energy changes during a reaction.
Reactant Energy: The level of energy before the reaction actually starts.
Product Energy: The level of energy after the reaction is completed.
Activation Energy: The energy necessary to start the reaction, indicated by the “hill” in the diagram.
Energy Change (ΔH): The difference in energy between the reactants and products. It tells you whether the reaction is exothermic or endothermic.
For exothermic reactions, the product energy is lower than the reactant energy, meaning energy is released.
For endothermic reactions, the product energy is higher than the reactant energy, meaning energy is absorbed.
3. Reversibility of Chemical Changes:
In chemistry, the reversibility of a chemical change is whether the reaction may be reversed so that the original substances are regenerated. Most chemical changes are virtually irreversible; however, some are reversible under certain conditions.
i. Irreversible Chemical Changes:
In many chemical reactions, the products formed are new substances which are difficult or impossible to change back into the original reactants by simple physical methods. These reactions are therefore irreversible.
Some examples of irreversibility include the following:
a. Combustion: Combustion of wood or fuel gives off carbon dioxide and water. It cannot easily be reversed.
b. Rusting of Iron: The iron reacts with oxygen and moisture to form iron oxide, known as rust, and it is not easy to reverse this without extreme conditions.
c. Cooking an Egg: If an egg is cooked, the proteins undergo irreversible chemical changes, so there is no way to get it back to its original raw state.
ii. Reversible Chemical Changes:
Reversible chemical reactions have products that can react with each other to produce the original reactants under certain conditions. Such reactions occur in both directions: from reactants to the products and vice versa, from the products forward to reactants.
These reactions are generally in dynamic equilibrium; the forward reaction rate equals that of the reverse reaction. This would imply that amounts of reactants and amounts of products are constant.
Example of Reversible Reaction:
This is the Haber process for the synthesis of ammonia. Ammonia (NH₃) can break down into nitrogen and hydrogen, and nitrogen and hydrogen can recombine to form ammonia. This reaction can go both ways under appropriate conditions of temperature and pressure.
iii. Factors That Influence the Reversibility of Reactions:
a. Temperature: Many reactions are reversible at a particular temperature or vice versa. Ammonia is one of them. Its production, the Haber process, is very sensitive to temperature and pressure.
b. Pressure: A change in pressure is enough to change the reaction by shifting the equilibrium for those involving gases.
c. Concentration: The concentration of reactants or products can sometimes be altered to favor the reaction towards the reactants or the products (Le Chatelier’s Principle).
4. Law of Conservation of Mass:
The Law of Conservation of Mass is one of the very basic principles of chemistry. The law states that mass is neither created nor destroyed in a chemical reaction. What this means is that the total mass of reactants before a reaction must equal the total mass of products after a reaction.
i. Key Points of the Law:
Mass Conservation: During a chemical reaction, atoms are rearranged but not created or destroyed. The total mass of the reactants will always equal the total mass of the products.
Closed System: This law is valid in a closed system where no mass is allowed to enter or leave during the reaction. If the system is open (for example, gases escaping), mass may appear to be “lost,” but it has just moved outside the system.
Example:
- Reaction:
- 4 grams of hydrogen (2 moles) combine with 32 grams of oxygen (1 mole) to form 36 grams of water (2 moles of H₂O). The mass before and after the reaction is the same, confirming the law of conservation of mass.
Implication in Chemical Equations:
The law states that chemical equations should be balanced. In a balanced equation, the number of atoms in each element on both sides of the equation will be equal so that no mass is lost or acquired.
ii. Examples of Law of Conservation of Mass
a. Burning of Wood: When wood burns, it reacts with oxygen to produce carbon dioxide, water, and ash. If you measure the mass of the wood and oxygen before the reaction and the mass of the products (carbon dioxide, water, and ash) after the reaction, the total mass remains constant. The apparent “loss” of mass is just the gases (CO₂, H₂O) being released into the atmosphere.
b. Rusting of Iron: Iron rusts if moistened and oxidized in the air. It combines to form iron oxide, also known as rust. The mass of iron and oxygen before rusting equals the mass of rust after rusting.
c. Chemical Engineering Applications:The conservation of mass is of utmost importance in industries like chemical engineering and pharmaceuticals. Here, very accurate reactions and product yields are required. This means that chemists can predict how much product will be formed based on the amount of reactants.