ACE IH Chemistry Physical Properties of Substances Review

Page 2: States of MatterOverview: This section discusses the three primary states of matter: solids, liquids, and gases. Each state has unique characteristics that define its behavior and properties.

• Example of Solid: Ice

• Example of Liquid: Water

• Example of Gas: Oxygen

Page 3: Kinetic Molecular TheoryKinetic Energy:

• Definition: Kinetic energy is the energy of motion. It is the energy an object has due to its movement.

• Relationship: As temperature increases, the kinetic energy of particles also increases. This means that particles move faster at higher temperatures. For example, in a hot gas, particles collide with more energy than in a colder gas.

• Example: A moving car has higher kinetic energy compared to a parked car; the faster it moves, the more kinetic energy it has.

Kinetic Molecular Theory:

• Concept: This theory states that matter is made of tiny particles (atoms and molecules) that are always in motion. The speed and type of motion vary depending on the state of matter.

• Behavior of Gas Particles: In gases, particles move quickly and randomly. They collide with each other and with the walls of their container. These collisions do not lose energy—this is called elastic collision.

• Example: Think of air inside a balloon; the gas particles inside are moving around quickly and pushing on the balloon walls, making it expand.

• Important Note: Gases don't have strong attractive forces between their particles, allowing them to spread out and fill their containers completely.

Standard Temperature and Pressure (STP):

• Definition: STP is a standard reference point used in science, defined as a temperature of 0 degrees Celsius and a pressure of 100 kPa (atmospheric pressure at sea level).

• Example: At STP, one mole of an ideal gas occupies approximately 22.4 liters, which helps in predicting gas behaviors in various conditions.

Gas Pressure:

• Origin: Gas pressure is caused by the collisions of gas molecules with the surfaces around them. More collisions lead to increased pressure.

• Example: In a soda can, gas pressure builds up because the carbon dioxide gas molecules collide with the can's interior walls.

Types of Motion:

• Vibration: In solids, particles vibrate in place due to strong attractions keeping them together.

  • Example: The molecules in a solid metal vibrate but stay in fixed positions and do not move around.

• Rotation: In liquids and gases, particles can spin around their own axes or rotate.

  • Example: Water molecules can rotate as they move, causing currents and waves.

• Translation: This refers to particles moving from one location to another, a behavior observed in liquids and gases.

  • Example: When you mix food coloring into water, the dye molecules move throughout the liquid to spread evenly.

State-specific Motion:

• Solids: Particles in solids are tightly packed together and can only vibrate in place. This gives solids a definite shape and volume.

  • Example: A cube of ice will maintain its shape and volume until it melts.

• Liquids: Particles have more energy and can slide past each other, allowing liquids to flow. They have a definite volume but take the shape of their container.

  • Example: A glass filled with water conforms to the shape of the glass while maintaining its volume.

• Gases: Particles in gases move freely and rapidly in all directions, allowing gases to completely fill their containers without a defined shape or volume.

  • Example: When you open a bottle of perfume, the scent quickly spreads throughout the room due to the fast-moving gas particles.

Page 4: Types of SolidsMolecular Solids:

• Characteristics: Made of molecules linked by weak intermolecular forces such as hydrogen bonds or van der Waals forces.

• Properties: Tend to be soft and have low melting points. Common in organic compounds like sugars.

• Example: Sugar (sucrose) is an example of a molecular solid that dissolves easily in water.

Ionic Solids:

• Composition: Formed from positively and negatively charged ions that are held together by strong ionic bonds.

• Properties: Have high melting and boiling points and are brittle; they conduct electricity when melted or dissolved in water.

• Example: Sodium chloride (table salt) is a typical ionic solid with a cubic lattice structure.

Metallic Solids:

• Bonds: Consist of metal atoms packed closely together, with a sea of delocalized electrons allowing for conductivity.

• Characteristics: Metals are generally malleable (can be hammered into shapes) and ductile (can be drawn into wires).

• Example: Copper is a metallic solid that's widely used for electrical wires due to its excellent conductivity.

Covalent Network Solids:

• Bonds: Composed of a vast network of atoms connected by strong covalent bonds.

• Properties: Extremely hard with high melting points and distinctive properties fitting their structure.

• Example: Diamonds are a classic example of a covalent network solid known for their hardness.

Page 5: Crystal StructuresCrystal:

• Definition: A solid material with a highly ordered and repeating 3D arrangement of atoms or molecules, often resulting in distinct geometric shapes.

• Example: Quartz crystals are known for their hexagonal form and clarity.

Amorphous Solid:

• Characteristic: Lacks a long-range order in arrangement, often appearing more like a liquid under certain conditions.

• Example: Glass is an amorphous solid, and it breaks irregularly rather than along clear crystal planes.

• Cleavage: Amorphous solids break unevenly, while crystalline solids break along flat planes.

• Fracture: Amorphous solids display irregular breakage, while crystalline solids have defined shapes upon breaking.

Unit Cell:

• Definition: The smallest repetitive unit that retains the overall structure and symmetry of the crystal.

• Example: In a face-centered cubic (FCC) structure, the unit cell is formed by atoms positioned at each corner and the center of each face.

Crystal Systems:

• Crystals can be categorized into different systems based on their unit cell symmetry:

  • Example of Cubic System: Sodium chloride (NaCl) exhibits a cubic crystal structure.

  • Example of Tetragonal System: Tin (Sn), an interesting metal that has a different arrangement than cubic metals.

  • Example of Orthorhombic System: Sulfur (S) can crystallize in this system, showing different angles between its axes.

Page 6: Phase ChangesOverview: This section explains how materials change between the three states of matter (solid, liquid, gas) and the energy changes that accompany these transitions.

Page 7: Types of Phase Changes

• Vaporization: When a liquid changes into a gas. This can occur through boiling at a specific temperature or by evaporation at lower temperatures.

  • Example of Boiling: Water boils at 100 degrees Celsius at sea level.

  • Example of Evaporation: Clothes dry on a line as water evaporates from their fabric, even on cool days.

• Condensation: The process of gas turning back into a liquid, often seen when water vapor collects on cooler surfaces.

  • Example: Dew forms on grass in the morning when dew points reach air saturation.

• Sublimation: The direct transition of a solid changing into a gas, skipping the liquid state.

  • Example: Dry ice (solid carbon dioxide) sublimates directly into carbon dioxide gas in room temperature air.

• Deposition: The reverse of sublimation, where a gas changes directly into a solid without becoming a liquid first.

  • Example: Frost forms on cold surfaces when water vapor in the air converts directly into ice crystals.

• Evaporation: Surface molecules of a liquid gain sufficient energy to break free and become gas.

  • Example: The fragrance from a perfume bottle fills a room as some molecules evaporate.

• Vapor Pressure: The pressure of gas molecules in equilibrium with its liquid phase in a closed container.

  • Example: The vapor pressure of water increases as the temperature rises, which influences boiling.

• Boiling Point: The temperature at which a liquid's vapor pressure equals the external atmospheric pressure.

  • Example: Water boils at lower temperatures in high-altitude areas due to lower atmospheric pressure.

• Impact of Altitude: At high altitudes, the decreased atmospheric pressure leads to reduced boiling points, which can affect cooking times.

Page 8: Phase DiagramsPhase Diagram:

• A graphical representation showing the different states of matter as conditions of temperature and pressure change.

  • Example: The phase diagram of water displays regions for solid, liquid, and gas, along with the conditions for melting and boiling.

• Triple Point: The specific conditions under which all three phases (solid, liquid, gas) coexist in equilibrium.

  • Example: Water’s triple point occurs at 0.01 degrees Celsius and 611.657 Pascal.

• Critical Point: The highest temperature and pressure a substance can reach, after which its liquid and gas phases cannot be differentiated.

  • Example: The critical point for water is 374 degrees Celsius and 22.06 MPa, beyond which water becomes a supercritical fluid.

Page 9: Ionic and Molecular CompoundsOverview: An analysis of ionic and molecular compounds, highlighting their unique characteristics and applications.

Page 10: Ionic vs. Molecular CompoundsIonic Compounds:

• Composition: Formed from metal cations (positively charged) and nonmetal anions (negatively charged) arranged in a regular lattice structure.

• Properties: Demonstrate high melting and boiling points, are usually solid at room temperature, and can conduct electricity when melted or dissolved in water.

• Example: Magnesium oxide (MgO) is an ionic compound known for its high melting point of 2852 degrees Celsius.

Molecular Compounds:

• Composition: Created through the sharing of electrons between nonmetal atoms (covalent bonding).

• Properties: Generally have lower melting and boiling points compared to ionic compounds and often exist as gases or liquids at room temperature.

• Example: Carbon dioxide (CO₂) is a simple molecular compound that exists as a gas at room temperature.

• Allotropes: Different structural forms of the same element are important for understanding unique properties.

  • Example of Allotropes of Carbon: Diamond (hard and clear), graphite (soft and slippery), and graphene (single layer of carbon atoms used in electronics).

Page 11: Comparing Metals and NonmetalsOverview: This section explores the key differences in properties and typical reactions of metals and nonmetals.

Page 12: Properties of Metals and Ionic Compounds

• Ductility: A measure of how much a material can be stretched into wires; metals like copper are very ductile, while ionic compounds tend to be brittle and break easily.

  • Example: Copper wires are drawn into thin strands for electrical applications.

• Malleability: Metals can be hammered or rolled into thin sheets without breaking or shattering.

  • Example: Aluminum foil can be flattened into thin sheets for cooking and food storage.

• Conductivity: Metals are exceptional conductors of heat and electricity due to the free movement of electrons; ionic compounds can only conduct electricity when dissolved.

  • Example: Gold is used in electrical connectors and circuits due to its excellent conductivity.

• Luster: Metals have a shiny appearance from their ability to reflect light, making them appealing for decorative purposes.

  • Example: Silver and gold are commonly used in jewelry for their luster.

• Alloys: Mixtures of different metals designed to enhance certain properties like strength or corrosion resistance.

  • Example: Steel is an alloy of iron and carbon that is commonly used due to its strength and durability.

Defects in Metals:

• Types: Defects in the arrangement of atoms can significantly affect the properties of metals.

• Interstitial Defect: Happens when an extra atom occupies a space between regular atoms, distorting the metal's structure.

  • Example: Carbon atoms in iron can strengthen steel through interstitial defects.

• Substitutional Defect: Occurs when one atom in the lattice is replaced by a different type of atom.

  • Example: Brass is made by substituting zinc for some copper atoms.

• Vacancy Defect: Refers to missing atoms in the crystal structure, which can impact the mechanical properties of the material.

  • Example: Vacancy defects in metals can enhance their strength by preventing slip during deformation.

• Dislocation: The presence of dislocations helps metals to deform easily during processes like forging or shaping, essential in manufacturing.

  • Example: Engineering metals such as lead can be easily shaped due to their dislocations.

Page 13: Metal Packing Arrangements

• Body-Centered Cubic (BCC): A structure where each atom is surrounded by eight others, usually resulting in lower ductility compared to other packing arrangements.

  • Example: Iron often adopts a BCC structure.

• Face-Centered Cubic (FCC): A configuration where atoms are more tightly packed with a coordination number of 12, enhancing ductility and malleability.

  • Example: Gold and copper typically crystallize in the FCC structure.

• Hexagonal Close-Packed (HCP): Each atom has twelve nearest neighbors, generally making materials with this structure less ductile.

  • Example: Zinc and cobalt typically exhibit HCP arrangements.

Page 14: Water and Aqueous SystemsOverview: This section provides a thorough examination of the special characteristics of water and how it interacts as a solvent with other substances.

Page 15: Water Properties

• Surface Tension: The cohesive forces between water molecules create a high surface tension, allowing some insects to walk on water surfaces.

  • Example: Water striders can skim across the surface of ponds due to its high surface tension.

• Solvation: Involves solvent molecules surrounding solute molecules to facilitate dissolving and chemical reactions.

  • Example: When salt dissolves in water, water molecules surround and separate the sodium and chloride ions.

• Density: Water’s density decreases when it freezes, which allows ice to float, providing insulation for aquatic life below.

  • Example: Icebergs float in oceans, allowing marine organisms to thrive beneath.

• Aqueous Solutions: Solutions where water is the solvent tend to dissolve polar substances well, while nonpolar substances do not dissolve.

  • Example: Sugar dissolves readily in water, but oil does not due to polarity differences.

• Electrolytes: Substances that dissolve in water to form ions, allowing the solution to conduct electricity, while non-electrolytes do not dissociate.

  • Example: Sodium chloride (table salt) dissociates in water, enabling it to conduct electricity.

Page 16: Properties of Solutions

• Dissolution Rate: Refers to how quickly a solute dissolves in a solvent, influenced by factors like temperature and surface area.

  • Example: Stirring a powder into water increases its rate of dissolution compared to simply letting it sit still.

• Solubility: The maximum amount of solute that can dissolve in a solvent at a specific temperature, typically expressed in grams per liter (g/L).

  • Example: At room temperature, a saturated solution of salt in water can hold about 357 g of NaCl per liter.

• Saturated Solution: A solution that has reached the maximum capacity of solute it can hold at a given temperature, where no more solute can dissolve.

  • Example: Adding salt to water until it no longer dissolves, with some settling at the bottom, signifies a saturated solution.

• Unsaturated Solution: A solution containing less solute than the maximum possible amount at that temperature.

  • Example: Water with just a small amount of sugar added can still dissolve more, indicating it's unsaturated.

• Supersaturated Solution: A solution that holds more solute than it normally would at equilibrium, making it unstable.

  • Example: When a concentrated sugar solution is heated and then cooled, it may remain liquid but will crystallize if disturbed.

• Solubility Curve: A graph showing how the solubility of a substance changes with temperature, providing insight into how dissolving behavior changes.

  • Example: The solubility curve for potassium nitrate illustrates that solubility increases significantly with temperature, providing crucial information for chemical processes.

Page 17: Study TipsGood luck studying!Reminder: Do not rely solely on this slideshow for exam preparation.Suggested Study Methods:

• Review previous ACE materials regarding bonding types and intermolecular forces for a stronger foundation.

• Regularly consult your class slideshows and notes for detailed information relevant to your topics.

• Complete worksheets and practice problems to consolidate your knowledge and identify areas needing improvement.

• Consult the Savvas textbook for extensive coverage of chemical concepts.

• Practice with questions from the Savvas question bank to gauge your understanding and application of the material as you prepare for exams.