GEN CHEM 2 PDF

**Q3: 2ND SEM** **TOPIC OVERVIEW** ------------------------------------------------------------------------------------------------------------------------------------ Good luck girl, you can do it. I know you can, look behind you. You have a lot of achievements, don\'t let your negativity ea...

**Q3: 2ND SEM** **TOPIC OVERVIEW** ------------------------------------------------------------------------------------------------------------------------------------ Good luck girl, you can do it. I know you can, look behind you. You have a lot of achievements, don\'t let your negativity eat you ------- --------------------------- **A** **Polarity of Molecules** ------- --------------------------- **Polarity** Equal or unequal sharing of electrons among atoms of molecules +-----------------------------------------------------------------------+ | **Polar Molecule** | | | | **There is unequal or asymmetrical** | | | | **distribution of electrons among the** | | | | **atoms of a molecule.** | +-----------------------------------------------------------------------+ - **OCTET RULE** *states that atoms tend to gain, share or* *transfer electrons in order to attain a* *stable 8 valance electron* ----------------------- **ELECTRONEGATIVITY** ----------------------- *Relative ability of an atom to draw electrons in a bond toward itself.* *Dipole or two poles. Oxygen is more electronegative than hydrogen* **Electronegativity differenc e-** *The greater the electronegativity difference, the greater the polarity of the molecule*. ------------------------ **Molecular Geometry** ------------------------ +-----------------------------------+-----------------------------------+ | **Molecular geometry** | | +===================================+===================================+ | **Linear** | Shape: If the molecule is linear, | | | it is nonpolar. | | | | | | 2 bond pairs | | | | | | No lone pairs | +-----------------------------------+-----------------------------------+ | **Bent or Angular** | 2 bond pairs | | | | | | 1 lone pair | +-----------------------------------+-----------------------------------+ | **Trigonal Planar** | 3 bond pairs | | | | | | No lone pair | +-----------------------------------+-----------------------------------+ | **Trihedral** | 4 bond pairs | | | | | | No lone pair | +-----------------------------------+-----------------------------------+ | **Trigonal Pyramidal** | 3 bond pairs | | | | | | 1 lone pair | +-----------------------------------+-----------------------------------+ | **Bent** | 2 bond pairs | | | | | | 2 lone pairs | +-----------------------------------+-----------------------------------+ | **Trigonal bipyramidal** | 5 bond pairs | | | | | | No lone pair | +-----------------------------------+-----------------------------------+ **NONPOLAR MOLECULE** The shape of the molecule is symmetrical - - - NONPOLAR MOLECULE-there is equal or symmetrical distribution of electrons among atoms in a molecule. ------- ------------------------------- **A** **THE INTERMOLECULAR FORCES** ------- ------------------------------- +-----------------------------------+-----------------------------------+ | **POLAR** | NONPOLAR | | | | | **-H-F,O,N** | -All diatomic molecules with same | | | element | | **-The molecule has lone pairs of | | | electron in its central atom** | -Molecules only contain C, H. | | | | | **-EN difference more than 0.5** | -The molecule has symmetry | | | | | | -EN difference is less than 0.5 | +-----------------------------------+-----------------------------------+ **INTERMOLECULAR FORCES** are the forces of attraction or repulsion between molecules, which influence the physical properties of substances, such as boiling and melting points, viscosity, and Solubility. These forces are generally weaker than intramolecular forces (chemical bonds within molecules) but play a significant role in determining a substance\'s behavior in different states of matter. are attractive forces between molecules that exert more influence in the condensed phase of matter- solid and liquid, and the non-ideal behavior of gases +-----------------------------------+-----------------------------------+ | **Types of Intermolecular | | | Forces** | | +===================================+===================================+ | **London Dispersion Forces** | -are weak attractions\ | | | that are used to explain the | | | attraction between\ | | | two nonpolar molecules due to | | | temporary\ | | | dipoles created by electrons' | | | motion. | | | | | | -by Fritz-London (German-American | | | physicist) | | | | | | The more electrons a molecule | | | has, the stronger the LDF. | | | | | | ex. Br2, has more electrons than | | | Cl2, so Br2 will stronger LDF | | | than Cl2, resulting in a higher | | | boiling point 59 degree celsius, | | | compared to -35degree celsius. | | | | | | London dispersion forces are | | | observed in nonpolar molecules. | | | These include: | | | | | | Halogens: | | | | | | Noble gases | | | | | | - | +-----------------------------------+-----------------------------------+ | **Dipole-Dipole Force** | -Dipole- Dipole Force | | | interactions occur when partial | | | charge form within a molecule | | | because of the uneven | | | distribution of electrons. | | | | | | -Polar molecules align so that | | | the positive end of one molecule | | | interacts with the negative end | | | of another molecule. | +-----------------------------------+-----------------------------------+ | **Ion-Dipole Forces** | -These forces exist when polar | | | molecules are attracted to ions. | | | It results when an ion and the | | | partial charge found at the end | | | of a polar molecule attract each | | | other. | | | | | | Attractive forces between an ion | | | and a polar molecule | | | | | | -Hydrogen Bond is a weak bond | | | formed when a hydrogen with | | | partial positive charge is close | | | to an atom in a molecule with | | | lone pairs of electrons or with | | | excess electron charge | | | (negative). | | | | | | -A Hydrogen Bond is an attractive | | | interaction between a hydrogen | | | atom bonded to an electronegative | | | F, O, or N atom and an unshared | | | electron pair of another nearby | | | electronegative atom. | +-----------------------------------+-----------------------------------+ +-----------------------------------------------------------------------+ | ***SUMMARY*** | | | | - - - - - - - | +-----------------------------------------------------------------------+ ------- --------------------------- **B** **Properties of Liquids** ------- --------------------------- -------- --------------------------------------------------- **B1** **Liquids due to intermolecular forces exhibits** -------- --------------------------------------------------- **VISCOSITY** *[Viscosity is the ability of a fluid to resist flowing motion.]* It is related to the movement of the molecules in the liquid and the IMF present. **What is the IMF? Intermolecular forces** Nonpolar molecules like CCl4 experience weak IMF tend to have low viscosity. [More polar molecules such as sugar syrup have high viscosity because of the H-bond among their -OH groups.] **VISCOSITY AND TEMPERATURE** The viscosity of a liquid decreases with increasing temperature because at higher temperature, the average KE of molecules that overcomes the attractive forces between molecules is greater. **SURFACE TENSION** Surface tension is the tendency of fluid surfaces to shrink into the minimum surface area possible. Molecules in the interior are attracted equally on all sides, whereas those at the surface are attracted only below and to the sides, producing a net inward force This inward force makes the molecules at the surface to pack closely together, causing the liquid to occupy the smallest area and behave like a tight skin. Like viscosity, [surface tension is higher in liquids that have higher IMF.] Both properties are temperature dependent, because at higher temperatures, molecules have more KE to counteract the attractive forces holding them *(intermolecular forces acting on a molecule in the surface layer of a liquid and the interior region of the liquid)* **SURFACTANTS** [Surfactants] like detergents, soap, and biological fat emulsifiers **decrease the surface tension of water and destroy the H-bonds.** **CAPILLARITY** This refers to the **spontaneous rising of a liquid in a narrow tube.** This [results from the cohesive forces (IMF) within the liquid] and [the adhesive forces between the liquid and the walls of the container.] When water is placed in a test tube, the surface or meniscus takes a ***U-shape (concave)*** ***af \> cf*** because adhesive forces are greater than cohesive forces. In mercury, the meniscus ***curved downward (convex)*** because the cohesive forces (metallic bonding)are stronger than the adhesive force. ***cf \> af*** +-----------------------------------+-----------------------------------+ | **Cohesion** | Adhesion | | | | | **refers to the attraction of | refers to the attraction of | | molecules for other molecules of | molecules with other molecules of | | the same kind, and water | the different kind | | molecules have strong cohesive | | | forces due to their H-bonds.** | | +-----------------------------------+-----------------------------------+ +-----------------------------------------------------------------------+ | **How does a tree get water from the ground all the way up to its | | leaves?** | | | | A tree gets water from the ground to its leaves through a process | | called capillary action, specifically involving cohesion, adhesion, | | and transpiration pull: [water molecules stick together | | (cohesion]) and to [the walls of xylem vessels | | (adhesion)], while water evaporating from the leaves | | (transpiration) creates a negative pressure that pulls more water | | upward from the roots. | +-----------------------------------------------------------------------+ **EVAPORATION** When liquid particles at the surface of [a liquid have enough kinetic energy to overcome the attractive forces of neighboring molecules,] they escape into gaseous or vapor state. This process is called evaporation. **EVAPORATION, KE, AND TEMPERATURE** The [escape of more energetic molecules on a liquid surface reduces the average KE of the remaining molecules.] Since KE is directly proportional to temperature, the liquid temperature decreases. **VAPOR PRESSURE** Vapor Pressure is the equilibrium pressure of a vapor above its liquid or solid state. **Volatility** [is the ability of the substance to vaporize.] *The higher the vapor pressure, the more volatile the liquid is.* *Molecules in vapor phase collide with the walls and lid of the container causing pressure.* *Increasing temperature increases the rate of evaporation and increases vapor pressure.* **BOILING POINT** The boiling point of a liquid is attained when the vapor pressure of the liquid matches the atmospheric pressure. That is, [as air pressure rises, the boiling point of the liquid rises,] and as atmospheric pressure falls, the boiling point of the liquid falls. -------- ------------------------------ **B2** **WATER AND ITS PROPERTIES** -------- ------------------------------ **WATER AS SOLVENT** Water molecules are attracted to other polar molecules and to ions. A charged or *polar substance that interacts with and dissolves in water* is said to be **hydrophilic.** In contrast, *nonpolar molecules like oils and fats do not interact well with water.* They are **hydrophobic.** The dissolving ability of water is the reason why it is considered as an important agent in the erosion of weathered materials. Wherever water flows, either through ground or bodies of water, it takes along valuable chemicals, minerals, and nutrients **Density** The mass per unit volume of a substance +-----------------------------------+-----------------------------------+ | **Density of a liquid water** | Density of ice | | | | | **1 g/cm3** | 0.9268 g/cm3 | +-----------------------------------+-----------------------------------+ The *water's lower density in its solid form is **due to the way hydrogen bonds are oriented as it freezes.*** +-----------------------------------------------------------------------+ | ***SUMMARY*** | | | | ***VISCOSITY*** | | | | *A measure of a fluid's resistance to flow.* | | | | ***SURFACE TENSION*** | | | | *It is the elastic force on the surface of a liquid. It is the amount | | of energy required to increase or stretch the surface of a liquid by | | a unit area.* | | | | *(intermolecular forces acting on a molecule in the surface layer of | | a liquid and the interior region of the liquid)* | | | | ***VAPOR PRESSURE*** | | | | *The pressure exerted by the vapor above the surface of the liquid in | | a closed container. It is the equilibrium pressure of a vapor above | | its liquid.* | | | | ***BOILING POINT*** | | | | *The temperature at which a liquid boils. Normal boiling point is the | | boiling point of a liquid when the external pressure is 1 atm.* | +-----------------------------------------------------------------------+ ------- -------------------------- **C** **THE NATURE OF SOLIDS** ------- -------------------------- **INTRODUCTION** A solid is formed when *the temperature of a liquid is low and the pressure is sufficiently high* *[causing the particles to come very close to each other.]* The general property of solids reflect *[the orderly arrangement of their particles and the fixed location of their particles.]* - - - When you heat a solid, its particles vibrate more rapidly as their kinetic energy increases. Expansion or increase in volume does take place, but this volume change is relatively small. ----------------------- **THERMAL EXPANSION** ----------------------- **Thermal expansion** is the *[process in which an object or body expands on the application of heat.]* Since the atoms are tightly packed in solids, thermal expansion is seen evidently here. **Melting Point** The melting point (mp) *[is the temperature at which a solid changes to a liquid.]* At this temperature, *[the disruptive behavior of particles are strong enough to overcome the attractions that hold them in fixed positions. ]* **Freezing Point** The freezing point (fp) *is the temperature at which a liquid changes to a solid.* **The melting and freezing point of a substance is of the same temperature.** At that temperature, the liquid and solid phases are at **equilibrium**. When a solid-liquid equilibrium is obtained, the melting point of the solid is equal to the freezing point in its liquid state. The temperature will not change as long as both phases are present. +-----------------------------------------------------------------------+ | Explanation: | | | | When a substance reaches its melting/freezing point, it enters a | | state of **phase equilibrium**, where the rates of melting (solid to | | liquid) and freezing (liquid to solid) are equal. At this | | temperature, [both phases coexist without any change in temperature | | because the system is **absorbing or releasing heat to change | | phase,** not to increase or decrease the temperature.] | | This latent heat of fusion ensures that the temperature remains | | constant until the phase transition is complete. | +-----------------------------------------------------------------------+ -------- ----------------------- **C1** **CLASSES OF SOLIDS** -------- ----------------------- **AMORPHOUS SOLID** **\"Without form\"** Solids whose particles **do not have orderly structures, they have poorly-defined shapes.** **Ex.** Glass, rubber, plastic etc These *solids are the result of melting, cooling and solidification of liquids before particles achieve internal order.* A **glass** is a transparent fusion *[product of inorganic substances that have cooled to a rigid state without crystallizing.]* **CRYSTALLINE SOLID** Solids **whose atoms, ions, or molecules are ordered in well-defined arrangements.** **Ex.** Salt, diamond, iron They are *[orderly arranged in a 3D pattern called Crystal lattice]* with uniform intermolecular force. +-----------------------------------+-----------------------------------+ | **TYPES OF CRYSTALLINE SOLIDS** | | +===================================+===================================+ | **ionic** | *Ionic crystals have ions as | | | constituent particles (cation and | | | anion)* | | | | | | - | | | | | | *Ex. Typical salts like NaCl, | | | MgCl2, Ca(NO3)* | +-----------------------------------+-----------------------------------+ | **Molecular** | *Molecular Crystals are those | | | which have molecules as | | | **constituent particles** | | | (particles that cannot be broken | | | down into smaller pieces at a | | | certain energy scale) as well as | | | structure units. [Weak | | | Intermolecular Force of | | | attraction holds them | | | together], low | | | melting point.* | | | | | | *Ex. HCl melts at -112 degree | | | celsius, frozen neon* | +-----------------------------------+-----------------------------------+ | **Metallic** | *Metallic Crystals- **are the | | | simplest type of structure** | | | since single metallic atoms are | | | constituent units.* | | | | | | *Ex. All metallic elements like | | | Cu, Na, Zn, Fe, Al* | +-----------------------------------+-----------------------------------+ | **Covalent Network** | *Covalent network- Atoms are | | | connected in a network of | | | covalent molecules.* | | | | | | *Ex. Diamond, Quartz, and Sugar* | +-----------------------------------+-----------------------------------+ -------- ----------------------------------- **C1** **CRYSTAL LATTICE AND UNIT CELL** -------- ----------------------------------- +-----------------------------------------------------------------------+ | **What determines the shape of a crystal?** | | | | **In a crystal, the particles are arranged in an orderly, repeating, | | 3D pattern called crystal lattice.** | +-----------------------------------------------------------------------+  The *[smallest portion of the crystal] which shows the complete pattern of its particles is called a [**unit cell**.]* When [unit cells are repeated in all directions, a crystal lattice is formed.] A **crystal lattice** is [a repeating array of any one of fourteen kinds of unit cells.] +-----------------------------------+-----------------------------------+ | **7 BASIC TYPES OF UNIT CELL** | | +===================================+===================================+ | **isometric** | **All three axes are equal** in | | | length, and all are perpendicular | | **(cubic)** | to one another. | +-----------------------------------+-----------------------------------+ | **Tetragonal | **Two of the three axes are equal | | ** | in length,** and all three axes | | | are perpendicular to another. | +-----------------------------------+-----------------------------------+ | **Orthorhombic** | **All three axes are unequal** in | | | length, and *[all are | | | perpendicular to one | | | another.]* | +-----------------------------------+-----------------------------------+ | **Hexagonal | Of four axes, three are of equal | | ** | length, are separated by equal | | | angles, and lie in the same | | | plane. **The fourth axis is | | | perpendicular to the plane of the | | | other three axes.** Hexagonal | | | cells have lattice points in each | | | of the two six-sided faces. | +-----------------------------------+-----------------------------------+ | **Triclinic** | **All three axes are unequal** in | | | length, and *[none is | | | perpendicular to | | | another.]* | +-----------------------------------+-----------------------------------+ | **Monoclinic | **All three axes are unequa**l in | | ** | length, and *[two axes are | | | perpendicular to each | | | other.]* | +-----------------------------------+-----------------------------------+ | **Rhombohedral (or trigonal)** | **All three axes are of equal | | | length,** and *none of the axes | | | is perpendicular to another*, | | | [but t*he crystal faces all have | | | the same size and | | | shape.*] | +-----------------------------------+-----------------------------------+ When sugar (such as sucrose) crystallizes, it typically forms a monoclinic unit cell. The structure depends on the specific type of sugar, but for sucrose, **3 TYPES OF CUBIC UNIT CELL** ------------------------------------------------------------- --------------------------------------------------------------------------------------- ------------------------------------------------------------------------------ **Simple cubic** has (8) atoms at each of the eight corners A **body-centered** **cubic unit cell** has additional atom in the center of its cube A **face-centered cubic cell** has additional atoms on each of its six faces ------------------------------------------------------------- --------------------------------------------------------------------------------------- ------------------------------------------------------------------------------ **CRYSTAL LATTICE AND UNIT CELL** The shape of a crystal reflects the arrangement of the particles within the solid. - -------- ---------------- **C1** **ALLOTROPES** -------- ---------------- The property of some chemical elements to exist in two or more different forms, in the same physical state, known as allotropes of the elements Some substances can exist in more than one form. +-----------------------------------+-----------------------------------+ | **Diamond** | - - | +===================================+===================================+ | **Graphite** | - - | +-----------------------------------+-----------------------------------+ | **buckminsterfullerene.** | - - | +-----------------------------------+-----------------------------------+ **Allotropes** The physical properties of diamond, graphite,and fullerenes are quite different. - - - **Only few elements have allotropes.** - ------- ------------------ **D** **PHASE CHANGE** ------- ------------------ **What comes to mind when you hear the word change?** **PHASE CHANGE** Phase change depends on several factors such as the nature of the substance, and the temperature and pressure within its environment. Phase change is always accompanied by a change in the energy of a system, classified into two: endothermic and exothermic. +-----------------------------------+-----------------------------------+ | **Endothermic** absorption of | **Exothermic** | | heat by the substance from the | | | environment. | Release of the heat by substance | | | to the environment | +-----------------------------------+-----------------------------------+ **EVAPORATION OR VAPORIZATION** A process by which a liquid is transformed into vapor. Due to the increased kinetic energy attraction force of its neighboring particles was overcome, most energetic particles escaped and entered the vapor phase or gas phase. **CONDENSATION** The reverse process of evaporation. It happens when particles in the vapor phase are cooled. Their kinetic energies are lowered, so the molecules move slower and begin to condense returning back to liquid phase. **SUBLIMATION** A solid passes directly to gas phase. Substances that sublime have relatively weak attraction force between the particles of the solid. **DEPOSITION** The change from gas to solid phase without passing the liquid phase, the reverse of sublimation. **REVERSE PHASE CHANGE** Reverse Phase Changes occur at the same temperature, and if they occur at the same rate, the phases are said to persist in dynamic equilibrium.  **HEATING CURVE** \(1) As heat is added, the temperature rises. \(2) After sometime, even with addition of heat, the temperature remains constant. At this point, a phase change is occurring. \(3) At what temperature does phase change occur? \(4) When phase change completed, temperature of the liquid rises again as the average KE is increasing, then a point is reached when added heat does not change the temperature; phase chang is occurring. (5)At the completion of the phase change, temperature again rises. \(6) How much heat has been added before the second phase change has began. \(7) What temperature was reached when 200J of heat was added? **SPECIFIC HEAT**  The amount of heat needed to raise the temperature of one gram of a substance by one Degree celsius. The following data can be used in calculations of heat transfer for water: When a material with small specific heat value absorbs energy, its temperature rises rapidly. That\'s why cooking pans are made of aluminum or copper. Aluminum and copper heat quickly and transfer the heat to the food being cooked. In contrast, materials with high specific heat values absorb a large amount of heat without much increase in temperature. **HEAT FUSION** The amount of energy required to overcome the intermolecular forces to convert a solid to liquid is Heat of Fusion. Unit is J/g- Joules per gram cal/g-calories per gram The amount of heat absorbed to melt one gram of ice to one gram of water is the heat of fusion of ice. Heat of fusion (0c)= 333.6 J/g(ice) Heat of fusion Q=m\^Hf **HEAT OF SOLIDIFICATION** the amount of heat release when one gram of water changes to one gram of ice. Q= m\^Hs **HEAT OF VAPORIZATION** The amount of energy necessary to convert a liquid into a gas is called heat of vaporization The amount of heat needed to raise the temperature of one gram of water by one Degree celsius. Q= m\^Hv **HEAT OF CONDENSATION** The amount of heat released when 1 gram of gas condenses to liquid. The heat of vaporization and condensation of a substance are equal but opposite in signs. Q=mc\^Hc ------- --------------- **E** **SOLUTIONS** ------- --------------- **NATURE OF SOLUTIONS** - - - - - Isoprophyl alcohol + water = miscible Oil + water = immiscible - 10mL of vinegar with 15mL of water, vinegar is the solute, and water is the solvent. - - - - Sodium chloride solution is an electrolyte. - Sugar solution is a nonelectrolyte +-----------------------------------------------------------------------+ | **Did you know?** | | | | Vitamin C is soluble in water. | | | | Vitamin A, D and E are soluble in fat. | | | | What vitamins are retained in the body? | | | | What vitamins are removed from the body with the urine? | +-----------------------------------------------------------------------+ ------------------------ **TYPES OF SOLUTIONS** ------------------------ **GASEOUS SOLUTIONS** The air we breathe is an example of a gaseous solution. Air has the properties of both nitrogen and oxygen. Trace amounts of argon and neon are also present in air. **LIQUID SOLUTIONS** Liquid solutions are the most common type of solutions. In liquid solutions, the solvent is always a liquid. The solute may be a gas, a liquid, or a solid. Carbonated beverages or soft drinks are solutions of carbon dioxide and other components in water. **SOLID SOLUTIONS** Gold or sterling silver jewelry are examples of solid solutions. The most common solid solutions are combinations of two or more metals called alloys. Copper + Tin = Bronze **Solute** **Solvent** **Solution** Example ------------ ------------- -------------- ---------------------------------- **Gas** **Gas** **Gas** Air (N2, O2, and other gases) **Gas** **Liquid** **Liquid** Carbonated Drinks (Co2 in water) **Liquid** **Liquid** **Liquid** Vinegar (acetic acid in water) **Liquid** **Solid** **Solid** Amalgam (mercury in silver) **Solid** **Solid** Solid Bronze **Solid** **Liquid** Liquid Sugar in water -------------------------------- **CONCENTRATION OF SOLUTIONS** -------------------------------- Solutions can be described qualitatively or quantitatively based on the amount of solute relative to a given amount of solvent. Qualitatively, we can describe solutions as either dilute or concentrated. A dilute solution contains a relatively small amount of solute, whereas a concentrated solution contains a relatively large amount of solute. For example, a solution containing 5 g of salt in 100 mL of water is dilute, while a solution containing 40 g of salt in the same volume of water is concentrated. ----------------------- **PARTS PER MILLION** ----------------------- Parts per million (ppm) is a unit for expressing very dilute concentrations. It is commonly used to express the concentration of pollutants in air or in water. Components of gas mixtures present in very small amounts are usually expressed in parts per million by volume as defined by the equation. If the solution is given in volume units, then you can simply change the mass units in the equation to volume units. ---------------------------- **MASS OR VOLUME PERCENT** ---------------------------- Percent (%) by mass of the component One of the ways to express concentration is by using mass. If the amount of the component and the solution are given in terms of volume, change the mass units to volume units. The volume percent is useful when dealing with liquid solutions. The volumes are usually given in milliliters. --------------------------------------- **MOLE FRACTION, MOLARITY, MOLALITY** --------------------------------------- **Mole Fraction (X)** is a way of describing solution composition. It is the ratio of the number of moles of one component of a mixture to the total number of moles of all components. The sum of the mole fractions of the components of a solution must be equal to 1. ***Xa** denoted as the mole fraction of the solute* ***Xb** denoted as the mole fraction of the solvent* **Molarity *(M)*** is the most common way of expressing the concentration of a solution. It is defined as the number of moles of solute per liter solution. **Molality *(m)*** of a solution is the number of moles of solute dissolved per kilogram of solvent. **Molarity *(M)*** is defined in terms of the volume of solution, while Molality *(m)* is defined in terms of the mass of a solvent. **MOLE** The mole is the unit of measurement in the International System of Units (SI) for amount of substance. It is defined as the amount of a chemical substance that contains as many elementary entities (e.g., atoms, molecules, or particles). This number is expressed by the Avogadro\'s constant, which has a value of 6.02 X 10\^23. To determine the number in a given mass Concentration [**Concentration of a solution** is a measure of the amount of solute that has been dissolved in a given amount of solvent or solution](https://www.canva.com/link?target=https%3A%2F%2Fchem.libretexts.org%2FBookshelves%2FIntroductory_Chemistry%2FBook%25253A_Chemistry_for_Allied_Health_%28Soult%29%2F08%25253A_Properties_of_Solutions%2F8.01%25253A_Concentrations_of_Solutions&design=DAF-pevMCDs&accessRole=viewer&linkSource=document). The more solute there is per amount of solvent, the more concentrated the solution gets. **Dilution** Dilution is the process of *["lowering the concentration of a solute in a solution by simply adding more solvent to the solution, such as water."]* Diluting a solution entails adding more solvent without adding more solute. +-----------------------------------------------------------------------+ | **Note** | | | | The number of moles of a solute does not change when a solution is | | diluted | +-----------------------------------------------------------------------+ Taking a higher concentration solution and adding water until the required concentration is reached is a popular way of creating a solution of a specific concentration. Dilution is the term for this procedure. Dilution can also be accomplished by **mixing a higher-concentration solution with a lower-concentration solution.** Since the total number of moles of solute does not change, ***M1V1=M2V2*** Where M1 and V1 are the molarity and volume of the initial solution, and M2 and V2 are the molarity and volume of the final solution, respectively. **Neutralizalization** An acid reacting with a base produces salt and water. This reaction is called neutralization. Neutralization can be carried out in the lab using titrimetric method or more commonly called **titration.** ***An acid or base*** of known concentration **(titrant)** is added in a stepwise increment to an acid or a base of unknown concentration (analyte). The point when the amount of hydronium ions (H+) is equal to the amount of hydroxide ions (OH-) is called ***the equivalence point*** or the ***endpoint. It can easily be determined using a pH meter.*** If an indicator is used, it is when the indicator changes color with the addition of the slightest amount of titrant. In conclusion, the amount of titrant balances the amount of analyte present during the reaction. **Solubility** Solubility refers to the maximum quantity of solute that can be dissolved in a given quantity of solvent at a given temperature. **Unsaturated Solution** Can still dissolve more solute An unsaturated solution has less solute than the solvent is capable of dissolving and the solute can dissolve completely. **Saturated solution** Holds the Maximum amount of solute at a given temperature In a saturated solution, there is enough solute present that if more were added, it would not dissolve and become a supersaturated solution. **Supersaturated Solution** A supersaturated solution contains more than the maximum amount of dissolved solute than a solvent is able to dissolve at a given temperature. **Recrystallization** Extra solute from crystals ---------------------------------- **FACTORS AFFECTING SOLUBILITY** ---------------------------------- 1. "Like dissolves like" Ionic and polar solutes are soluble in polar solvents, while nonpolar solutes are soluble in nonpolar solvents. **Effect of the Nature of Solute And solvent on Solubility** -------------------------------------------------------------- ------------------- ---------------------- **Solute** **Polar Solvent** **NonPolar Solvent** ionic Soluble insoluble polar soluble insoluble nonpolar insoluble soluble 2. The solubility of most molecular and ionic solids increases with temperature. Exception: Sodium sulfate which increase its solubility in decreasing temperature. For gases, all gases become less soluble as temperature increases. Rather dissolve more at lower temperature. 3. According to William Henry (an English chemist), the solubility of a gas at a given temperature is directly proportional to the partial pressure of the gas over the solution. This is known as the Henry's Law. The solubility of gases in liquids is greatly affected by pressure. It is directly proportional to pressure, increasing pressure increases solubility of the gas in water.