Results for "Solvent"

Common Organic Solvents

Protic and Polar Aprotic Solvents

chapter 7 (solvents)

SN1/SN2/E1/E2 + Protic/Aprotic Solvents

Chapter 4 The Effects of Chemical Reactions. • Introduction to Chemical Reactions. - Chemical reaction: a process in which one or more substances change into one or more new substances. - Clues that a chemical reaction has occurred : 1. Color change Example: two colorless aqueous solutions mix together to produce a bright yellow precipitate. 2. A precipitate (solid) is formed when mixing two solutions together. 3. Gas formation. Bubbles of gas (effervescence) are produced when mixing substances together (solid – liquid or aqueous – aqueous ….) 4. Heat is produced. - Chemical reactions are described by using word equations or chemical equations. - Chemical equations need to be balanced when written because it shows the correct proportions (amounts) of chemicals in a reaction. - A balanced chemical equation has equal number of atoms of each element in the reactants (left hand side) and the products (right hand side). - Exercise: Balance the following equations. a) KClO3→KCl + O2 b) Na2O + H2O NaOH c) Cu + AgNO3 Cu(NO3)2 + Ag d) C3H7OH + O2 CO2 + H2O • Synthesis and Decomposition Reactions. Synthesis: Two or more substances (elements and / or compounds) combine to form one larger compound. General pattern: A + B → C Examples: N2 + 3 H2 → 2 NH3 CaO + CO2 → CaCO3 2 P + 3 Cl2 → 2 PCl3 Decomposition: This is opposite to synthesis; that is, one large compound breaks down (decomposes) into 2 or more simpler substances. Example: 2 KClO3 → 2 KCl + 3 O2 General pattern: R → S + T Remark: Usually decomposition happens due to heat or electricity. - Predicting the product of decomposition or synthesis reactions. 2 AlCl3 (s) → 2 Al (s) + 3 Cl2 (g) Zn (s) + S (s) → ZnS (s) 2 Zn (s) + O2 (g) → 2 ZnO(s) - Single Displacement (Replacement) Reactions. Definition: A reaction in which an element displaces (replaces) another element in a compound, producing a new compound and a new element. General pattern: A + BC → AC + B Example: Mg (s) + CuSO4 (aq) → MgSO4 (aq) + Cu (s) Zn (s) + 2 AgNO3 (aq) → Zn(NO3)2 (aq) + 2 Ag (s) Fe (s) + MgCl2 (aq) → no reaction. Remark: The element that displaces the other element in a compound must be more reactive (active) than that element, otherwise no reaction takes place. In the general pattern above, A should be more reactive than B for the reaction to proceed. The following reactivity (activity) series lists the chemical strength (reactivity) of the metals in order from the more reactive to the less reactive. KPlease stop calling my amazing zebra in the long Nahungry class. sorry !! Ca Mg Al Zn Fe Sn Pb H Cu Ag Examples of single displacement reactions : 2 Al (s) + 3 CuSO4 (aq) → Al2(SO4)3 (aq) + 3 Cu (s) Sn (s) + Zn(NO3)2 (aq) → no reaction Exercise: Complete and balance the following equations. If there is no reaction occurring write no reaction. a) 2 Al (s) + 6 HCl (aq) → 2 AlCl3 (aq) + 3 H2 (g) b) Cu (s) + H2SO4 (aq) → no reaction c) 2 AlCl3 (aq) + 3 Ca (s) → 3 CaCl2 (aq) + 2 Al (s) d) Mg (s) + 2 HNO3 (aq) → Mg(NO3)2 (aq) + H2(g) - Reactivity of halogens decreases down the group. F2> Cl2> Br2> I2 The reactions taking place for the halogens or their compounds are in solution (aqueous) Examples: Cl2 (aq) + 2 KBr (aq) → 2 KCl (aq) + Br2 (l) Cl2 (aq) + NaF (aq) → no reaction. Exercise: F2 (aq) + 2 LiCl (aq) → 2 LiF (aq) + Cl2 (g) I2 (aq) + NaCl (aq) → no reaction • Double displacement reactions. - Definition: A reaction in which two compounds mix together and an exchange of ions (elements) occurs which results in the formation of 2 new compounds. - General pattern: AB + CD → AD + CB - Solubility: the amount of solute that dissolves in a given amount of solvent at a given temperature. - When we say a substance is soluble, it means it dissolves in water; whereas if it is insoluble it means it doesn’t dissolve in water. - The compound in a reaction that is soluble is in aqueous (aq) phase, whereas the compound which is insoluble is in the solid state (s). - The solid which is formed in a double displacement reaction is called the precipitate and it is insoluble. - Solubility rules (used in double displacement reactions). 1. All alkali metal ions and ammonium ion (NH4+) are soluble. 2. All nitrates (NO3-) are soluble. 3. All sulfates (SO4-2) are solubleexceptwith Ba+2 , Pb+2 , Ca+2 , Sr+2 , Ag+ . 4. All chlorides, bromides and iodides(Cl-, Br-, I-) aresolubleexcept with Ag+ , Pb+2 , Hg+, Cu+ 5. All OH- are insolubleexceptwith rule 1, and Ba+2 and Sr+2 . 6. All oxides (O2-), sulfides (S2-), sulfites (SO32-), carbonates (CO32-), phosphates (PO43-) are insoluble except with rule 1 Remark: If all compounds formed in a double displacement reaction are soluble (aqueous) then no reaction takes place. Exercise: State whether each of the following compounds is soluble or insoluble ? Na2SO4 : Fe(NO3)2: LiOH: ZnSO4: PbBr2: BaSO4: Mg(OH)2: PbO: NH4Cl: Na2S: Cu(OH)2: KF: Exercise: Complete and balance the following chemical equations: - KNO3 (aq) + NaCl (aq) → - LiCl (aq) + AgNO3 (aq) → - Zn (s) + FeSO4 (aq) → - NaOH (aq) + CuCl2 (aq) → - ZnCl2 (aq) + Na3PO4 (aq) → - Pb(NO3)2 (aq) + K2S (aq) → • Net ionic equation: a chemical equation which shows ONLY the ions that are involved in the formation of the precipitate (solid). Examples: Pb+2 (aq) + S-2 (aq) → PbS (s) Ag+ (aq) + Cl- (aq) → AgCl (s) Cu+2 (aq) + 2 OH- (aq) → Cu(OH)2 (s) • Full ionic equation: an equation which shows All the ions in the soluble (aqueous compounds) in both reactants and products. Example: - 2 NaOH (aq) + CuCl2 (aq) → 2 NaCl (aq) + Cu(OH)2 (s) 2 Na+ (aq) + 2 OH- (aq) + Cu+2 (aq) + 2 Cl- (aq) → 2 Na+ (aq) + 2 Cl- (aq) + Cu(OH)2 (s) - 3 ZnCl2 (aq) + 2 Na3PO4 (aq) → Zn3(PO4)2 (s) + 6 NaCl (aq) Full ionicequation: 3 Zn+2(aq) + 6 Cl-(aq) + 6 Na+ (aq) + 2 PO4-3 (aq) → Zn3(PO4)2 (s) + 6 Na+ (aq) + 6 Cl- (aq) Net ionic equation: 3 Zn+2 (aq) + 2 PO4-3 (aq) → Zn3(PO4)2 (s) Exercise: Complete and balance the following equation, then write full ionic and net ionic equations for the reaction. Pb(NO3)2 (aq) + 2 NaI (aq) → Full ionic equation: Net ionic equation: Spectator ions: the ions that are not involved in the formation of the precipitate (solid). Note that the spectator ions appear on both sides of the full ionic equation. For example, in the above reaction, Na+ (sodium ions) and NO3- (nitrate ions) are the spectator ions. Exercise: Complete and balance the following equation, then write the net ionic equation and identify the spectator ions. BaCl2 (aq) + K2SO4 (aq) → Net ionic equation: Ba+2 (aq) + SO4-2 (aq) → Spectator ions: - Combustion reaction is a special type of (synthesis) reaction in which the substance reacts with (burns in) oxygen. Examples: C(s) + O2(g) → CO2(g) • Production of gases (lab scale): 1. CO2 2. SO2 3. H2 4. H2S (hydrogen sulfide) 5. NH3 (ammonia) General pattern of the chemical reactions to produce the above gases: 1. Metal carbonate + acid → CO2 Example: Na2CO3 (aq) + 2 HCl (aq) → 2 NaCl(aq) + CO2(g) + H2O(l) 2. Metal sulfite + acid → SO2 K2SO3 (aq) + 2 HCl (aq) → 2 KCl(aq) + SO2(g) + H2O(l) 3. Metal + acid → H2 Remark: This is a single displacement reaction therefore the metal used in the reaction should be higher in the reactivity series than hydrogen. Zn (s) + 2 HCl (aq) → ZnCl2 (aq) + H2(g) 4. Metal sulfide + acid → H2S Na2S (aq) + 2 HCl (aq) → 2 NaCl (aq) + H2S (g) 5. Ammonium compound + base (alkaline solution) → NH3 NH4Cl (aq) + NaOH (aq) → NaCl (aq) + NH3 (g) + H2O (l) Exercise: Write the net ionic equations for each of the above 5 reactions. Answers 1. 2 H+ (aq) + CO3-2(aq) → CO2(g) + H2O (l) 2. 2 H+ (aq) + SO3-2(aq) → SO2(g) + H2O (l) 3. Zn(s) + 2 H+(aq) → Zn+2(aq) + H2(g) 4. 2H+ (aq) + S-2 (aq) → H2S (g)

Water: The Solvent for Biochemical Reactions

Medical Toxicology 2: Pesticides, Metals, Solvents, and Food

ch. 6 solvents

Solvents

chapter 6: solute solvent

LECTURE 9 - Solvent extraction and Chromatography basics

Organic Chemistry Solvents

Polar Protic vs Aprotic Solvents

**Properties of Matter** - **Physical vs Chemical Properties** - *Physical Properties:* Can be observed without changing the substance (e.g., color, density, melting point, boiling point) - *Chemical Properties:* Describe a substance’s ability to undergo chemical changes (e.g., flammability, reactivity with acid) - **Examples of Each** - Physical: Ice melting, water boiling, density, solubility - Chemical: Rusting iron, burning wood, tarnishing silver - **Intensive vs Extensive Properties** - *Intensive:* Do not depend on the amount of matter (e.g., density, boiling point, color) - *Extensive:* Depend on the amount of matter (e.g., mass, volume, length) --- **Density** - **Definition, Units, and Formula** - Density (ρ) = Mass (m) / Volume (V) - Units: g/cm³ (solids), g/mL (liquids), kg/m³ (gases) - **Comparison of Densities** - Solids: Generally highest density - Liquids: Lower density than solids but higher than gases - Gases: Lowest density - **Solid: Regular vs Irregular Shape** - *Regular Shape:* Use geometric formulas to find volume - *Irregular Shape:* Use water displacement method - **Factors Affecting Density** - Temperature (increase decreases density for most substances) - Pressure (affects gases significantly) - Composition (different materials have different densities) --- **Elements, Compounds, and Mixtures** - **Matter: Definition and Examples** - Anything that has mass and takes up space (e.g., air, water, rocks) - **Pure vs Impure Matter** - *Pure:* Elements and compounds (e.g., oxygen, water) - *Impure:* Mixtures (e.g., saltwater, air) - **Atom vs Element** - *Atom:* Smallest unit of an element - *Element:* Substance made of one type of atom - **Compounds vs Mixtures** - *Compounds:* Chemically bonded elements (e.g., H2O, CO2) - *Mixtures:* Physically combined substances (e.g., salad, air) - **Types of Mixtures** - Homogeneous (solutions, uniform throughout) - Heterogeneous (distinct parts, not uniform) - **Examples of Mixtures** - Homogeneous: Saltwater, air - Heterogeneous: Salad, granite --- **Solubility** - **Solutions** - *Parts:* Solute (dissolved substance) + Solvent (dissolving substance) - *Examples:* Saltwater (solute: salt, solvent: water) - **Effect of Temperature and Pressure** - Higher temperature increases solubility of solids in liquids - Higher pressure increases solubility of gases in liquids - **Gases vs Liquids** - Gases dissolve better in cold liquids under high pressure - Liquids dissolve better at higher temperatures - **Gaseous, Liquid, and Solid Solutions** - Gaseous: Air (oxygen in nitrogen) - Liquid: Saltwater (NaCl in H2O) - Solid: Alloys (brass, steel) - **Concentrations** - Unsaturated: Can dissolve more solute - Saturated: Maximum solute dissolved - Supersaturated: Holds more than normally possible - **Solubility Curve** - Shows solubility vs temperature - Higher points indicate higher solubility --- **The Mole** - **Avogadro’s Number** - 6.022 x 10^23 particles per mole - **Molar Mass** - Mass of one mole of a substance (g/mol) - **Particle, Mass, and Mole Calculations** - Particle calculations: Using Avogadro’s number - Mass calculations: Converting between grams and moles - Mole calculations: Determining amount of substance - **Moles at STP (Standard Temperature and Pressure)** - 1 mole of gas = 22.4 L at STP - **Atomic Mass Units (AMU)** - Unit for atomic/molecular mass --- **Labs** - **Density Lab** - Measure mass and volume, calculate density - Compare densities of different materials - **Elements, Compounds, and Mixtures Lab** - Classify substances based on their properties - **Mystery Powder Lab** - Identify unknown substances using solubility and reactions - **Cornstarch Lab** - Explore properties of non-Newtonian fluids This guide covers essential concepts in matter, density, solubility, and the mole, along with relevant lab activities

aprotic solvents

Lab 4: Separating Acids and Neutral Compounds by Solvent Extraction

CHEM 50: WATER: SOLVENT FOR BIOCHEMISTRY

Gases, Vapors and Solvents

HOMEOSTASIS Maintaining a stable internal environment respond to stimuli Reacting to changes in the environment reproduce and develop Creating new organisms and growing adapt and evolve Changing over time to better suit the environment INDUCTIVE REASONING Making generalizations based on specific observations DEDUCTIVE REASONING Making specific predictions based on general principles Matter Anything that has mass and takes up space elements Substances that cannot be broken down into simpler substances protons Positively charged particles in the nucleus neutrons Neutral particles in the nucleus electrons Negatively charged particles orbiting the nucleus Atomic Number Number of protons in an atom Isotopes Atoms of the same element with different numbers of neutrons Octet Rule Atoms tend to gain, lose, or share electrons to achieve a full outer shell of 8 electrons molecule Two or more atoms held together by chemical bonds compound A substance consisting of two or more different elements IONIC BONDS Bonds formed by the transfer of electrons COVALENT BONDS Bonds formed by the sharing of electrons reactants Starting materials in a chemical reaction products Ending materials in a chemical reaction WATER solvent Dissolves many substances WATER cohesion & adhesion Water molecules stick to each other and other surfaces WATER high surface tension Water's surface resists being broken WATER high heat capacity Water can absorb a lot of heat without changing temperature WATER heat of vaporization Water requires a lot of energy to evaporate WATER varying density Ice is less dense than liquid water acidic solutions Solutions with a pH below 7 basic solutions Solutions with a pH above 7 pH scale Measures the acidity or basicity of a solution buffers Substances that resist changes in pH Organic Molecules Molecules containing carbon carbon The backbone of organic molecules functional groups Chemical groups attached to carbon that give molecules specific properties Macromolecules Large molecules made up of smaller subunits monomers The individual subunits of a polymer polymers Long chains of monomers Dehydration Synthesis Reaction Joins monomers by removing water Hydrolysis Reaction Breaks polymers by adding water Role of Enzymes Speed up chemical reactions Carbohydrates monosaccharides glucose Simple sugars Carbohydrates disaccharides glycosidic bonds Two monosaccharides joined together Carbohydrates polysaccharides starch glycogen cellulose Many monosaccharides joined together LIPIDS Glycerol & Fatty Acids saturated Fatty acids with no double bonds LIPIDS Glycerol & Fatty Acids unsaturated Fatty acids with double bonds PROTEINS Enzymes Proteins that catalyze chemical reactions PROTEINS amino acids peptide bonds The monomers of proteins, joined together PROTEINS protein structure primary The sequence of amino acids PROTEINS protein structure secondary Local folding patterns (e.g., alpha-helices and beta-sheets) PROTEINS protein structure tertiary The overall 3D shape of a single polypeptide PROTEINS protein structure quaternary The arrangement of multiple polypeptides in a protein conformation The 3D shape of a protein denaturation The unfolding of a protein DNA Deoxyribonucleic acid, the genetic material RNA Ribonucleic acid, involved in protein synthesis ATP Adenosine triphosphate, the energy currency of the cell

O Chem Common Solvents