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solution chemistry

This is a mind map about solution chemistry. The main contents include: Chapter two: Electrolyte Solution, Chapter three: Buffer Solution, Chapter one: Solution.

Edited at 2024-11-06 19:31:39

21_ethanSawyer

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solution chemistry

21_ethanSawyer

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Outline

Introduction to the book 'Solution Chemistry'
- 34
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Recommended online resources for 'Solution Chemistry'
- 16
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Laboratory safety regulations for 'Solution Chemistry'
- 48
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The Relationship between Solution Chemistry and Other Disciplines
- 19
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The Application of Solution Chemistry in Daily Life
- 23
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Experimental Techniques of Chromatography
- 17
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Experimental Techniques in Solution Chemistry
- 14
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Classification of Solution Chemistry
- 38
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Basic Overview of Solution Chemistry
- 19
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solution chemistry

Chapter one:Solution

1-1 Basic Terms of Solution

Solution

solvent

solute

mixed case

miscible

immiscible

Solution Process

diffusion

solvation (hydration)

Absorption and release of heat

endothermic

exothermic (exothermic)

1-2 Solubility of substance

Solubility & Factors that affect solubility

The nature of solutes and solvents

temperature

Stress has little effect

1-3 Concentration of Solution

1-3.1 Amount-of-substance and Molar Mass

Amount-of-substance

mole

The mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12.

one mole of anything contains 6.022 × 1023 entities. 6.022 × 1023 is called Avogadro's Number

molar mass

unit：g/mol

1.3.2 Amount-of-substance Concentration

Amount-of-substance Concentration (molarity)

Unit: mol/L or mmol/L……

mass of a substance (m) Amount of substance (n) = -------------------------- molar mass (M)

Unit: mole (mol)

Confusing errors

So, c(bB)=1/bc(B)

1-3.3 Reacting Rule of Equal Amount of Substance

When aA tT = dD eE

n(aA)= c(aA)V(A) n(tT)= c(tT)V(T)

1-3.4Mass Concentration (This is not density, this is the mass of the solute, and density is the mass of the solution)

Ρ=m/v

The SI unit is kg/m3 (equal to g/L)

1.3.5 Volume fraction

Va/Va Vb

1-3.6 Molality (mB )or(bB)

mB=nB/mA

kilogramsofsolvent(mA)

unit must be kg

amount-of-substanceofsolute(nB)

In Chinese textbooks, it is represented by b

1-3.7 Amount-of-substance Fraction(xi)(Mole fraction)

1-3.8Mass Percent of Solute

mass of solute/mass of solution

Summary

G1: molarity; mass concentration; Volume fraction

G2: molality; mass percent of solute; mole fraction

1-4 Collaborative Properties of Solutions

*1-4 Colligative Properties of Solutions

Colligative properties are properties that depend only on the number of solute particles in solution and NOT on the nature (identity) of the solute particles, regardless of whether they are atoms, ions, or molecules

Colligative Properties of NonelectrolyteSolutions (Nonvolatile, non-electrolyte and relatively dilute solutions).

Colligative Properties of Electrolyte Solutions.

1-4.1 Vapor Pressure Lowering of Solution -Raoult’sLaw

Raoult’s law : Vapor pressure lowing of a dilute solution that containing nonvolatile, nonelectrolyte is directly proportional to the amount-of-substance fraction of the solute and has nothing to do with the nature of solute.

Vapor Pressure of Solution (p): p= poxA

unit: Pa, kPa

po: vapor pressure of pure solvent

xA:the mole fraction of the solvent in the solution

For Solution: solvent (A), solute (B)

∵ xA xB =1, p= p°xA = p°(1- xB )= p°– p°xB ∴ p°- p = p°xB Δ p =p°- p = p°xB

For a dilute solution: mB = n B(bB=nB)

Δ p = K mB

K=p°MA /1000; mB : molality of B

p = px A Δp = px B Δp = K m B

1-4.2Elevation of Boiling Point

ΔTb= Tb-Tb°= K'Δp= K'K mB =Kb mB

Kb: the modal boiling-point elevation constant of solvent

1-4.3Depression of Freezing Point

Freezing point depression refers to the lowering of the freezing point of solvents upon the addition of solutes.

ΔTf= Tf°-Tf =KΔp = K'K mB=Kf mB

1-4.4Osmosis & Osmotic Pressure

Osmosis: the process of solvent flow through asemipermeable membrane from apure solvent or from a dilutesolution to a more concentrated solution in order to equalize theconcentrations of solutes on the two sides of the membrane.

osmotic pressure (π):The pressure of a solution which just stops osmosis

The relationship between osmotic pressure, concentration and temperature (Van’t Hoff Law)

Π= c RT

unit: Pa, kPa

For a dilute solution: Π= mB RT

because mB=cB,the density of solution is close to pure water

1-4.5 Using Colligative Properties to Determine Molar Mass

For dilute nonelectrolyte solutions:

Δp= Kb mB, ΔT b= Kb mB , ΔT f = K f mB , Π= c RT ≈ mB RT

For dilute electrolyte solutions:

Δp’ =i Kb mB, ΔT b’=i Kb mB ΔT f ’= i K f mB, Π’= i c RT ≈ i mB RT

i: van’t Hoff factor

Example: NaCl, KNO3, i =2; CaCl2, MgCl2 i = 3;

1-4.7Medical Applications of Osmosis

1-4.7.1 Isotonic, Hypotonic, Hypertonic Solution

◆Hypotonic: contains less solute compared to another solution.

◆Hypertonic: contains more solutes compared to another solution.

◆Isotonic: two solutions have the same concentration of solutes

1-4.7.2Osmolarity

It is the total concentration of ALLosmotically active solute particles in the solution.

Osmole (Osm), Cos , Osm/L, mOsm/L nonelectrolytes:OsmL=mol/L For electrolytes: Osm/L=mol/L×i

1-4.7.3Crystalloid Osmotic Pressure and Colloidal Osmotic Pressure

Crystalloid Osmotic Pressure

Definition:Small molecular crystal substance (NaCl (≈ 80%), NaHCO3, glucose) formed π

Physiological action:Regulate balance of fluid and electrolyte on the two sides of cell membrane

Semipermeable membranes:Cell membrane

Colloidal Osmotic Pressure

Definition:Large molecular colloid substance (plasma proteins) formed π

Physiological action:Regulate balance of fluid and electrolyte on the two sides of blood capillary wall

Semipermeable membranes:Capillary wall

1-4.7.4Applications of Osmosis

Dialysis

The Artificial Kidney

Reverse Osmosis

Chapter two:Electrolyte Solution

2-1 Strong and Weak Electrolyte Solution

2-1.1 Theory of Strong Electrolyte Solution

Ion-ion Interaction Theory

form ionic atmophere

The ions restrict each other, causing the ion movement rate to slow down, resulting in a decrease in the ionization degree measured by conductivity.

apparent degree of ionization

The degree of mutual restraint of ions in reactive strong electrolytes

The greater the apparent ionization degree, the smaller the restraint effect between ions.

The smaller the apparent ionization degree, the greater the containment effect between ions.

Ion Activity and Activity Coefficient

Actual concentration of ion (c) multiply a correction factor - activity coefficient ( f ).

a = f·c

Generally, a<c, 0<f<1

Activity coefficients are influenced by

ion concentration

the electric-charge number of ion

Ionic Strength ( I )

I =1/2（c₁z₁² c₂z₂²…cizi²)

Where, I is ionic strength;

c is the amount-of-substance concentration of the ion i;

z i is the charge number of the ion i

2-1.2 Ionization Equilibrium of Weak Electrolyte Solution

The law of Chemical Equilibrium(Equilibrium Constant)

a A b B c C d D

Ionization Constant (Ki )

Ka stands for “acid constant”

Kb stands for “base constant”

Degree of Ionization (α)(Percent Ionization)

1. Definition:

2. The factor influencing degree of ionization

① the nature of solute:

② the initial concentration of solute:

③ temperature

Dilution Law

(1) the weak electrolyte must be monoprotic

(2) α≤5%

The Common Ion Effect and Salt Effect

1. Common ion effect

The ionization of a weak electrolyte is markedly decreased by the addition of an ionic compound containing one of the ion of the weak electrolyte, this effect is called the common ion effect.

2.Salt effect

The ionization of a weak electrolyte is increased by the addition of an soluble strong electrolyte which not contains the common ion with the weak electrolyte. This effect is called salt effect.

2-2 Theory of Acid-base

2-2.1 The Arrhenius theory of acids and bases (Arrhenius ionization theory)

2-2.2 Bronsted-Lowry Acids and Bases (Acid-base Proton Theory)

1. Definition of acid and base

Conclusion:

Acid or base may be a molecule, atom, or ion.

Some molecules or ions are capable of donating a proton, and also accepting a proton, which named ampholyte.

There are no concepts of salt in acid-base proton theory.

2. Essence of Acid and Base Reaction

Essence: proton transfer reaction.

3. Relative Strength of Acids and Bases

① Judgment by acid-base reaction:

② Compare by K a or Kb

③ The relationship between acid-base strength and solvent (proton-accepting ability)

4. The Leveling and Differentiating Effect

The leveling effect:

The inability of a solvent to differentiate the relative strengths of all acids stronger than the solvent’s conjugate acid is known as the leveling

The differentiating effect:

The ability of a solvent to differentiate the relative strengths of all acids stronger than the solvent’s conjugate acid is known as the differentiating effect.

2-3 Acidity and Calculation of Solution

2-3.1 Automation of Water

2-3.2 Acidity of solution (The pH Function)

2-3.3 Calculation of Acidity of Solution

For Strong Acids and Bases

Monoprotic Weak Acids and Bases

K times the concentration under the square root (the same goes for acids and bases)

2-4 Equilibrium between Dissolution and Precipitation

2-4.1 Solubility Product Constant (Ksp)

2-4.2 Exchange between Solubility and Ksp

● AB (s , ksp)

● AB2(A2B)

2-4.3 Formation/dissolution of precipitation

Rule of Solubility Product

Formation of precipitation

condition: Qi>Ksp

Selective Precipitation of Ions (Separation of Ions by Fractional Precipitation)

Selective precipitation is a technique of separation ions in an aqueous solution by using a reagent that precipitates one or more of the ions, while leaving other ions in solution

Dissolution of precipitation

condition: Qi＜Ksp (ion product < solubility product)

methods to dissolve precipitation.

1. Forming weak electrolytes by adding some compounds make precipitation dissolve.

2. Forming coordination compounds by adding some agents make precipitation dissolve.

3. Producing oxidation-reduction reactions by adding oxidizing agents or reducing agents make precipitation dissolve.

Chapter three:Buffer Solution

3-1 Concept of Buffer Solution

3-1.1 Concept of Buffer Solution

Buffer solution (buffer): A solution that can resist changes in pH when limited amounts of strong acid or base are added to it.

● Adding a small amount of acid

● Adding a small amount of base.

3-1.2 Composition and Type of buffer solution

3-2 The pH of Buffer Solution

3-2.1 Henderson-Hasselbalch equation

buffer ratio

When the temperature of the same buffer solution is constant, the pH value depends on the buffer ratio

3-2.2 Calculating the pH of buffer solution

3-3 Capacity of Buffer Solution

3-3.1 Concept of buffer capacity

The buffer capacity (β) is the amount-of-substance of strong acid or base per liter needed to produce a unit change in pH.

Δb: the amount-of-substance of strong acid/base added

ΔpH: pH variation after addition of Δb mol strong acid/base

unit: (mol L-1pH-1)

3-3.2 Factors Influencing Buffer Capacity (β)

Table 3-1 the relationship between capacity and concentration

Table 3-2 The relationship between capacity and buffer ratio

3-4 Designing and Preparation of Buffers

3-4.1 Preparing Principle of Buffer Solution

1. Choose the suitable buffer system

2. Choose the best buffer system

3. Choose the suitable total concentration

Buffer capacity: 0.01-0.1

3-4.2 Methods of Preparing Buffer Solution

• Choose the suitable buffer system

• Choose the suitable total c

• Calculate the amount of acid and base

• Mix the amount of acid and base

3-4.2.1Ways to make a buffer

3-4.3 Preparing a Buffer in Lab

methods

Calculation and weighing

Dissolution

pH measurement

3-5 Buffer Systems in the Body

The pH of blood: 7.35~7.45

Maintaing Body pH Balance

Intracellular buffers: HHbO2-HbO2-, HHb-Hb-,

Intracellular buffers in RBC:HHbO2-HbO2-, HHb-Hb-, H2CO3-HCO3-

We can not determine their strength, because H3O is the strongest acid that can exist in aqueous solution.

Body fluids: about 300 mOsm/L. Isotonic solution: cos = 280～320mOsm/L Hypertonic solution: cos> 320mOsm/L Hypotonic solution: cos < 280mOsm/L

Hemolysis

Normal

Plasmolysis/creation

Significant Figures

Significant Figures Rules

General rule: 'significant figures are any non-zero digits or trapped zeros. They do not include leading or trailing zeros'

1.Trailing zeros in a whole number are not significant (unless they come from a measurement). Examples: 300, 300 kg

2.Leading zeros are zeros before non-zero numbers. Example: 00054

3.Trapped zeros are zeros placed in between any two non-zero digits. Example: 5001

4.Zeros placed between the dot and the first non-zero digit are not significant if the number is smaller than 1 (it is a particular case of the rule above). Example: 0.069 has 2 significant figures.

5.Trailing zeros at the end of the number after the decimal point are significant. Example: 812.00 has five significant figures.

6.Only zeros: The number 0 has one significant figure. Therefore, any zeros after the decimal point are also significant. Example: 0.00 has three significant figures, 0.000 has four significant figures and so on.