AP Chemistry 4.4 Physical and Chemical Changes Study Notes - New Syllabus Effective fall 2024
AP Chemistry 4.4 Physical and Chemical Changes Study Notes- New syllabus
AP Chemistry 4.4 Physical and Chemical Changes Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Explain the relationship between macroscopic characteristics and bond interactions for:
i. Chemical processes.
ii. Physical processes.
Key Concepts:
- Physical & Chemical Changes
- Representing Chemical Changes
- Balancing Chemical Equations
- Physical & Chemical Processes
4.4.A.1Chemical vs. Physical Processes:
1. Chemical vs. Physical Changes:
| Property | Chemical Change | Physical Change |
|---|---|---|
| Definition | Involves breaking and forming of chemical bonds, resulting in new substances. | Involves changes in the state or appearance of a substance, but no new substance is formed. |
| Molecular Changes | Atoms are rearranged to form new compounds. | No change in the molecular structure; only the form or state changes. |
| Type of Bonds Affected | Chemical bonds (strong bonds like covalent, ionic, etc.). | Intermolecular forces (weaker forces like hydrogen bonds, van der Waals forces, etc.). |
| Energy Changes | Significant energy changes (heat, light, electricity) are often involved. | Energy changes typically involve heat (e.g., melting or freezing) but not significant chemical energy changes. |
| Reversibility | Often irreversible or difficult to reverse. | Generally reversible (e.g., water can freeze and melt repeatedly). |
| Examples | – Burning wood (combustion) – Rusting iron – Digesting food | – Melting ice – Boiling water – Dissolving sugar in water |
| Formation of New Substances | Yes, new substances with different chemical properties are formed. | No, only changes in form or appearance occur without creating new substances. |
2. Chemical Processes:
| Type of Reaction | Energy Change | Example |
|---|---|---|
| Combustion | Exothermic (releases heat) | Burning methane or gasoline |
| Respiration | Exothermic (releases ATP energy) | Cellular respiration (glucose → ATP) |
| Photosynthesis | Endothermic (absorbs sunlight) | Plants converting CO₂ and H₂O into glucose |
| Neutralization | Exothermic (releases heat) | HCl + NaOH → NaCl + H₂O |
| Decomposition | Exothermic (releases heat) | Hydrogen peroxide breaking down |
| Ionic Bond Formation | Exothermic (releases heat) | Formation of NaCl |
| Haber Process | Exothermic (releases heat) | Formation of ammonia (NH₃) |
3. Physical Processes:
Changes of states in matter are called phase changes. A change in intermolecular forces between molecules may happen by either the addition or removal of energy, which is generally heat.
i. Important Intermolecular Forces:
a. London Dispersion Forces (Van der Waals forces)
b. Dipole-Dipole Interactions
c. Hydrogen Bonds
The state of a substance, that is, whether it is a solid, a liquid, or a gas, and how readily it would change phases is determined by these forces. Let us consider what occurs with intermolecular forces during some common phase changes.
a. Melting (Solid to Liquid):
Process: A solid absorbs heat and hence its molecules are able to vibrate more strongly, breaking the intermolecular forces which had held them into a fixed structure.
Role of Intermolecular Forces: Molecules in a solid are tightly bound into a rigid structure through strong intermolecular forces-in the case of ice, these are hydrogen bonds; in salts, ionic bonds.
When it melts, the applied heat wins over the intermolecular forces holding the molecules together so that they can slip past one another and become a liquid
For instance, water melting to ice
b. Freezing (Liquid to Solid)
Process: A liquid sheds its heat, and its molecules slow down to form structures or solids
Role of Intermolecular Forces: When the temperature drops, the kinetic energy of the molecules is low, and intermolecular forces (hydrogen bonds or van der Waals forces) may be able to hold the molecules together in a solid structure.
Example: Water becomes ice.
c. Boiling (Liquid to Gas)
Process: A liquid is heated to give enough energy to the molecules so they can get free from intermolecular forces and escape into the gaseous state.
Role of Intermolecular Forces: In a liquid, the molecules are held together by forces such as hydrogen bonding (for water) and van der Waals for non-polar liquids.
These forces can be broken when there is sufficient heating, causing the molecules to escape from the liquid as a gas.
Example: Water turning to steam.
d. Condensation (Gas to Liquid)
Process: There is loss of heat, and molecules lose their energy, making intermolecular forces significant to permit the transformation of gas into liquid.
Role of Intermolecular Forces:
Gases are of sufficient energy in order to resist intermolecular forces; thereby, these would not bond up.
As and when the gas molecules lose some energy, their intermolecular forces are found to be highly significant to pull the gases so that a liquid is formed.
Example: Water vapor directly changes into water vapour through sublimation.
e. Sublimation (Solid to Gas)
Process: A solid transforms directly into a gas bypassing the liquid phase
Role of Intermolecular Forces:
In the solid, the molecules are very strongly held together by strong intermolecular forces.
Sublimation happens when that much energy is supplied to thoroughly overcome these forces so that the molecules could straight escape into the gas phase.
Example: Dry ice (solid CO₂) subliming into CO₂ gas.
f. Deposition (Gas to Solid)
Process: A gas transforms directly into a solid without going through the liquid phase.
Role of Intermolecular Forces: Deposition is said to occur when the gas molecules lose energy, and the intermolecular forces between the molecules become strong enough to pull the gas molecules into a solid state bypassing the liquid phase.
Example: Frost formation on a cold surface (water vapor turning directly into ice).
Role of Intermolecular Forces During Phase Changes:
Stronger Intermolecular Forces: Those compounds that exhibit stronger forces between molecules (for example, the hydrogen bonds of water or the ionic bonds of salts) require more energy to change phase and, thus, have higher melting or boiling points.
Weaker Intermolecular Forces: Compounds that exhibit weaker forces between molecules, such as those of nonpolar molecules, change phase with relative ease at lower temperatures-for example, methane boils at much lower temperature than does water.
4. Energy Considerations:
i. Chemical Reactions-Changes in Energy:
Chemical reactions break and form bonds among atoms. There are two broad types of reaction types: exothermic and endothermic. These reaction types depend upon the nature of the bonds involved.
a. Exothermic Reactions:
In an exothermic process, the energy is given out during the process of a reaction.
Rationale: Since energy needed to break bonds is already less than what has to be imparted to the new bonds that get formed to gain energy in their formation.
Examples:
Combustion (for example, combustion of methane: CH₄ + O₂ → CO₂ + H₂O + energy)
Respiration: glucose + oxygen → carbon dioxide + water + energy
Products are less energetic compared to reactants. In the exothermic reaction, the energy released comes from the surroundings in the form of heat, light, or sound.
b. Endothermic Reactions:
Energy Change: Energy is absorbed from the surroundings.
Explanation: The amount of energy to break the bonds in the reactants is greater than the energy released when the new bonds are formed in the products.
Photosynthesis (6CO₂ + 6H₂O + light → C₆H₁₂O₆ + 6O₂ )
Dissolving salts in water
e.g. NaCl in H2O
In endothermic reactions the energy of products is higher as compared to energy of reactants. Energy absorbed from surroundings is in the form of heat
Balanced View of Energy
It absorbs energy to break up the bonds. The bond formation is exothermic because this is a process of energy discharge.
ii. Energy Changes relating to Phase Changes:
A phase transition is a change of state, from liquid to solid, solid to liquid, or even from a solid to a gas. It is driven by either the presence or absence of heat. No bonds are being made or broken in this process but intermolecular forces are rearranged.
a. Exothermic Phase Transitions:
As one phase transforms to a lower-energy phase from the higher-energy phase, energy is transferred from the surroundings into the system.
Examples
Liquids freezing into the solid state. A cooling liquid is transferring heat energy out of the liquid to the surroundings. Intermolecular forces grow stronger, and a substance may change state from liquid to solid.
Condensing gas to liquid. Cooling gas undergoes its molecules losing energy, then coming together in droplets. Energy is released as heat during both freezing and condensing.
b. Endothermic Phase Transitions:
Energy Change: A substance absorbs energy from the surroundings when it changes from a lower-energy phase to a higher-energy phase.
Examples:
Melting (solid to liquid): Absorbs heat because the energy gained by the solid molecules is such that it has enough strength to break intermolecular forces and enter the liquid state.
Boiling (liquid to gas): Heat absorbed is used in overcoming the forces between the intermolecular particles in the liquid, allowing those molecules to come out into the gas phase.
Energy absorbed as heat, both for melting and in boiling.
4.4.A.2 Breaking of Chemical Bonds: Physical or Chemical Process? :
1. Ionic Bonding and Dissolution:
The ionic bonds between the sodium ions, Na⁺ and the chloride ions, Cl⁻ in salt, usually sodium chloride, NaCl, are broken when the salt dissolves in water. The ions spread in the water. The process is accompanied by energy changes as well as interactions of the ions with the molecules of water.
i. The Ionic Bond in Salt (NaCl):
Ionic Bonds: Sodium chloride (NaCl) is composed of sodium that loses an electron and gives up to chlorine, and so there is a positively charged sodium ion (Na⁺) and a negatively charged chloride ion (Cl⁻). Opposites attract, so this makes for an extremely strong ionic bond in between the two ions.
This ionic bond is extremely strong holding ions together in a three-dimensional solid lattice.
ii. Dissolution Process:
Adding salt to water causes the following to happen:
a. Ion Breaking – Ionic Bond Separation:
Polarity must be involved here. The molecule is polar; therefore, the water molecule has a partial charge on the hydrogen atoms of a positive nature and on the oxygen atom of a negative nature.
Water molecules surrounding the Na⁺ ions and Cl⁻ ions in the salt have some of their oxygen atoms negatively charged (partially negative) that attract the positive Na⁺ ions and the hydrogen atoms positively charged (partially positive) attracting the negative Cl⁻ ions.
The force of attraction of the water molecules toward the ions is strong enough to break the ionic bond binding Na⁺ and Cl⁻ within the salt crystal.
b. Hydration of Ions:
As a result, the Na⁺ and Cl⁻ ions dissociate and get hydrated.
Each Na⁺ ion gets surrounded by water molecules with the oxygen atoms facing the positive Na⁺ ion.
Each Cl⁻ ion gets surrounded by water molecules with the hydrogen atoms facing the negative Cl⁻ ion.
c. Ion Dispersion:
Once the ions (Na⁺ and Cl⁻) are hydrated, they are free to roam about in the water, thus a sodium chloride solution in water.
This dispersion of the ions in the solvent allows the salt to dissolve, and an electrolyte solution which conducts electricity.
iii. Energy Considerations:
Breaking the Ionic Bond: Lattice energy is the energy required to break the ionic bond between Na⁺ and Cl⁻. It is liberated when the ions are surrounded by water molecules.
Hydration Energy: The energy liberated when the water molecules surround and stabilize the ions is hydration energy.
When the hydration energy liberated is greater than the lattice energy, then the dissolution process is energetically favorable and the salt dissolves.
iv. Overall Process:
The process of dissolving salt in water involves two major energy changes:
Breaking the ionic bond, which requires energy to break the lattice energy.
Hydration of ions, in which energy is released since ions are stabilized by water molecules.
For NaCl, hydration energy is typically greater than lattice energy, and thus the process of dissolving salt in water becomes an exothermic process; that is, it releases energy.
2. Physical vs. Chemical Process:
| Aspect | Physical Process (Dissolution) | Chemical Process |
|---|---|---|
| Nature of Change | No new substance is formed. | New substances are formed through the breaking/formation of chemical bonds. |
| Composition of Substances | The chemical composition of the solute (e.g., NaCl) does not change. | The chemical composition changes to form new substances. |
| Bonding | Only intermolecular forces (e.g., ion-dipole interactions) are involved; no new chemical bonds are formed. | Chemical bonds are broken and formed, resulting in new chemical species. |
| Reversibility | Typically reversible: The solute can be recovered by evaporating the solvent (e.g., salt can be recrystallized). | Often irreversible: New substances may not be easily recovered in their original form. |
| Example | Dissolving salt (NaCl) in water: NaCl dissociates into Na⁺ and Cl⁻ ions without changing its chemical identity. | Burning paper: Paper reacts with oxygen to form ash and gases, creating new compounds. |
| Energy Change | May involve energy absorption or release (e.g., endothermic or exothermic), but no chemical bonds are formed or broken. | Involves energy changes (e.g., heat released or absorbed) due to bond formation or breaking in a chemical reaction. |
| Chemical Identity of Solute | The solute (e.g., NaCl) retains its original chemical identity as ions in solution. | The solute undergoes a chemical change and becomes a new substance with a different chemical identity. |
3. Ion-Dipole Interactions:
Ion-dipole Interactions are one of the intermolecular forces that exist between an ion, and a polar molecule which is similar to the water. They are particularly significant in a process such as the dissolving of salts in water, where molecules of water encapsulate and stabilize the ion.
i. Formation of Ion-Dipole Interactions:
a. Water Molecule Structure :
Water (H₂O) is polar molecule as there is partial positive charge associated with the hydrogen atoms (δ⁺), and partial negative charge with oxygen atom (δ⁻).
Such a polarity causes uneven distribution of the charge among water molecules, thereby the molecules of water can have an interaction with ions.
b. Ions in Solution:
Ions are charged particles formed from the dissociation of ionic compounds (such as NaCl in a solvent. For example, sodium chloride (NaCl) is composed of sodium ions (Na⁺) and chloride ions (Cl⁻).
c. Ion-Dipole Interactions:
When ionic compounds like NaCl dissolve in water, the surrounding water molecules make ion-dipole interactions with each other.
For Na⁺ (positive ion): The oxygen atoms (partially negative) of the water molecules are attracted to the Na⁺ ions. The oxygen atoms point towards the Na⁺ because they are partially negative.
For Cl⁻ (negative ion): The hydrogen atoms (partially positive) of the water molecules are attracted to the Cl⁻ ions. The hydrogen atoms point towards the Cl⁻ because they are partially positive.
d. Stabilization of Ions:
These ion-dipole interactions stabilize the ions in solution by surrounding them with water molecules. This prevents the ions from re-associating with each other and keeps them dissociated in the solution.
ii. Energy Used to Form Ion-Dipole Forces:
Exothermic reaction: The energy released in the ion-dipole forces because the attraction between water and ions stabilizes the system, which is part of why many ionic compounds are exothermic in solution when dissolved in water.
Hydration energy: The energy released when water molecules surround and stabilize ions is called hydration energy. Hydration energy is an important portion of the energy in a dissolution and clearly will be used to counter the lattice energy of the ionic solid.
iii. Ion-Dipole vs. Other Intermolecular Forces:
Ion-Dipole forces are stronger than dipole-dipole and hydrogen bonding because the ion has a full charge, as opposed to partial charge on the dipoles.
Especially important for the solvent-solute interactions when trying to dissolve salts and other ionic compounds in polar solvents such as water.
Physical vs Chemical Changes (OLD CONTENT)
- A physical change is a change in the form of a substance, not in its chemical composition = reversible
- A physical change can be used to separate a mixture into pure compounds, but will not break compounds into elements
- Ex: phase changes, separation techniques (Distillation, filtration, chromatography), dissolving
- Involve changes in intermolecular forces
- Physical properties: shape, solubility, density, volume, & interparticle distance at diff temperatures (& phases) may change but composition stays the same
- A physical change can be used to separate a mixture into pure compounds, but will not break compounds into elements
- Chemical change: a substance becomes a new substance with different properties and composition
- Ex: rusting, bunsen burner (burning), acid, combustion (reacted with oxygen)
- In chemical changes, one substance changes to another by reorganizing the way the atoms are attached to each other → involve changes in intramolecular forces
- Bonds have been broken, and new ones have been formed
- Indicators of chem change: temp, light, or color change, gas or precipitate formation
Example: Dissolving NaCl
- Can be argued both ways bcuz is reversible (physical change) and ionic bonds are broken (chemical)