Copy of The Grind_ Treinta Días de Infierno Chem only.pdf

The Grind: Treinta Días de Infierno Follow us @s0grindbam4tf ! Notes are made by: @ovenedpeas Izyan (3E) @stevoltatingtreacheries Isaac (3D) Team @deoxyfication & @s0grindbam4tf wishes you a successful EOY 2024! Sec 3 Chemistry & Physics EOY Survival Guide Chemistry Experimental Chemistry (24/8) Kinetic Particle Theory (25/8) Atomic Structure (26/8) Chemical Bonding (26/8 & 27/8) Structure and Properties of Materials (27/8) Chemical Formulae and Equations (28/8)* Mole Concept and Stoichiometry (28/8)* Acids & Bases (31/8)* Salts (31/8)* Ammonia (31/8)* Oxidation and Reduction (31/8)* Chemical Energetics (31/8)* The Periodic Table (1/9)* The Reactivity Series (1/9)* Experimental Chemistry Apparatus & Measurements: Measurement of Time Digital Stopwatch (+/- 0.01s) Analogue Stopwatch (+/- 0.1s) Measurement of Temperature K = C+273 (Kelvin to degree Celsius) Measurement of Length Metre rule (+/- 0.1cm) Measuring tape (+/- 0.1cm to 0.5cm) 1m = 10dm = 100cm = 1 000mm Measurement of Mass 1kt = 1 000t 1t = 1 000kg 1kg = 1 000g = 1 000 000mg Measurement of Volume Pipette (accurate & fixed) - usually 20.0cm^3 or 25.0cm^3 Volumetric Flask (ACCURATE + FIXED) Measuring Cylinder (range of volumes, ≈ 0.5cm^3) 31.5cm^3, 23.0cm^3 Burette (range of volumes, ≈ 0.05cm^3) - 31.55cm^3, 23.00cm^3 1m^3 = 1 000dm^3 = 1 000 000cm^3 Gas syringe (measures up to 100cm^3) Reading Measurements: Reading volumes of liquid Position your eyes at the MENISCUS (the convex/concave curve formed at the surface of the liquid. For the maths kiddos; the turning point of the surface of the liquid) to avoid PARALLAX ERROR When measuring mass, ensure that measuring apparatus are not considered in the mass to prevent ZERO ERROR Collection of Gases: The method used to collect gas depends on these 2 following physical properties of the gas Solubility in water Density (take Mr of air to be 30) When gas is insoluble in water: Water Displacement When gas Mr > 30, denser than air: Downward Delivery When gas Mr < 30, less dense than air: Upward Delivery When gas Mr = 30, equally dense to air: Downward/Upward delivery Drying of Gases: Background information: Sometimes, the gas collected needs to be dried before used in experiments to prevent any impurity. Hence, we dry the gases Concentrated Sulfuric Acid: Dry acidic/neutral gases NO ALKALINE If gas (alkaline; ammonia) can react with it, HELL NO. If gas (acidic/neutral; chlorine) cannot react with it, HELL YES Quicklime/Calcium Oxide: Dry alkaline/neutral gases NO ACIDIC Absorbs carbon dioxide from air Must be freshly heated to use Cannot dry gases reacting with calcium oxide (e.g. carbon dioxide) Dry gases which are alkali (e.g. ammonia) Fused Calcium Chloride: Dry any gas but AMMONIA Readily absorbs moisture from the air Must be freshly heated to use Cannot dry ammonia Separation of Substances in Mixtures: Solid-Solid Mixtures: Magnetic Attraction (magnetic materials S.I.N.C. vs. other materials) If one solid component in the mixture is MAGNETIC, then it can be easily separated from the NON-MAGNETIC solids A magnet is used Sieving (big particle vs. small particle) If the mixture consists of BIGGER and SMALLER particles, they can be separated by using a SIEVE with suitable PORE SIZE A sieve is used Using Suitable Solvents (soluble vs. insoluble) If one of the solids is soluble in a particular solvent, then that particular solvent can be used to separate the solid from the mixture Sublimation (sublimates vs. does not sublimate) If one of the substances sublimes (changes from SOLID to GASEOUS state DIRECTLY) on heating while the other substances are stable at sublimation temperature, then sublimation can be used Ammonium Chloride, Iodine and Naphthalene are substances that can sublime Gas can be collected by scraping via providing a cool surface for it to change its state back into a solid, known as the sublimate Pure sublimate can be scraped off the cool surface for collection Solid-Liquid Mixtures: Filtration (Factors of Sieving + Solubility) Used to separate INSOLUBLE SOLIDs from LIQUIDS Filter funnel + Filter paper is used Filtrate: Liquid that passes through the filter paper Residue: Solid that remains on the filter paper Evaporation to Dryness (Solvent evaporated, leave behind dissolved solid) Used to separate a DISSOLVED SOLID from its SOLVENT by HEATING the mixture until all the solvent has EVAPORATED Substance with lower B.P will turn into a gas first, leaving the other substance behind. A SIGNIFICANT DIFFERENCE in B.P between both substances is needed to separate the dissolved solute from its solvent. Crystallisation (Evaporation to Dryness, but solvent is saturated) *Saturated = no more solid can dissolve in solvent* Used to obtain a pure solid from its SATURATED SOLUTION It is a “gentler” method than evaporation to dryness and is used How to carry out Crystallisation: 1. GENTLY heat the solution in an EVAPORATING DISH until the solution is saturated 2. COOL the solution gradually until crystals appear within the solution 3. Carefully pour the mixture through a funnel lined with filter paper to collect the crystals 4. Wash the crystals with COLD DISTILLED WATER to remove impurities 5. Dry between sheets of filter paper Simple Distillation (Pure solvent evaporated from solution & condensed) Used to separate a pure solvent (LIQUID) from a solution Quite literally the opposite of Evaporation to Dryness Liquid-Liquid Mixtures Miscible: mixed together, uniform/homogenous solution Immiscible: separated into layers overtime, heterogeneous solution by density HOMOgeneous: can mix HETEROgenous: cannot mix, immiscible liquids separate into layers called phases Separating Funnel (separate immiscible liquids, use of bottom tap) Used to separate immiscible liquids Chromatography (substances in mixture have different solubilities) Used to separate a mixture of substances which have DIFFERING SOLUBILITIES in a GIVEN SOLVENT Start line marked with pencil - graphite insoluble in solvent Solvent front below start line - prevent dyes dissolving into solvent Small dotting of original dye - prevent OVERLAPPING in separation Edge of chromatogram submerged in solvent Solvent absorbed by paper, travels to other edge, separating dyes Further distance travelled = higher solubility in solvent and Rf *Rf value = distance travelled by substance/solvent < 1 Fractional Distillation (Distillation, substances have similar BP) Includes use of fractionating column & boiling chips (-> smooth) Heating solution causes both substances to evaporate and rise Substance with higher BP condenses on cool beads in the fractionating column, returns into flask Substance with lower BP exits column into condenser for condensation and collection as pure substance Remove conical flask with pure substance after BP of other substance is reached to prevent contamination Liquids of similar BP require longer fractionating column Purity of substances: - Pure substance has fixed MP & BP - Mixtures/Impure substance has MP & BP over a range - Salt added to water (raise BP & lower FP) to prevent ice formation Kinetic Particle Theory Kinetic Particle Theory: All matter is made up of tiny particles and these particles are in constant random motion. Adding/removing energy from particles affects forces experienced by them, determining their physical states The Differences between Solid, Liquid & Gaseous Particles: State of Matter Solids Liquids Gases Energy needed to A lot Lesser Little break arrangement Particle VERY CLOSELY CLOSELY PACKED in a VERY FAR APART in Arrangement * PACKED in ORDERLY DISORDERLY manner DISORDERLY manner manner Attractive Forces Very strong Less strong Very weak B/W Particles Kinetic Energy of Very low Low High Particles Particle Movement * VIBRATE/ROTATE SLIDE past one Move QUICKLY/ about FIXED another FREELY RANDOMLY in ALL POSITIONS throughout liquid directions Shape Definite Indefinite Indefinite Volume Definite Definite Indefinite Compressibility No No Yes Conversion by Heat Gain: AS SOLID HEATS UP: 1. THERMAL ENERGY is converted to KINETIC ENERGY of the particles [ENERGY CONVERSION] 2. Particles VIBRATE AND ROTATE faster about their FIXED positions [PARTICLE MOVEMENT] 3. The TEMPERATURE of the substance rises towards (MELTING POINT) [CONCLUSION] AS SOLID IS MELTING: 1. THERMAL ENERGY is ABSORBED from the surroundings and the TEMPERATURE of the SOLID is at MELTING POINT [ENERGY ABSORPTION + MELTING POINT] 2. The particles with INCREASED ENERGY can OVERCOME the INTERMOLECULAR FORCES OF ATTRACTION in the SOLID state [OVERCOMING OF IFOA] 3. The ORDERLY PACKING arrangement of the particles is DISRUPTED and they start to slide past one another freely [DISRUPTION OF ARRANGEMENT + PARTICLE MOVEMENT] 4. BOTH SOLID and LIQUID is present during the MELTING process [PRESENCE OF SOLID + LIQUID] 5. The TEMPERATURE remains CONSTANT throughout the MELTING process until ALL the substance has MELTED [CONCLUSION] AS LIQUID HEATS UP: 1. After ALL SOLID has melted, THERMAL ENERGY is converted to KINETIC ENERGY of the particles [ENERGY CONVERSION] 2. Particles SLIDE PAST one another with INCREASING SPEED [PARTICLE MOVEMENT] 3. The TEMPERATURE of the liquid RISES towards (BOILING POINT) [CONCLUSION] AS LIQUID EVAPORATES: 1. THERMAL ENERGY is ABSORBED from the surroundings, and the TEMPERATURE of the liquid is at BOILING POINT [ENERGY + TEMPERATURE] 2. The particles with INCREASED ENERGY can OVERCOME the INTERMOLECULAR FORCES OF ATTRACTION in the LIQUID state [OVERCOMING OF IFOA] 3. The particles move FURTHER APART, QUICKLY and RANDOMLY [DISRUPTION OF ARRANGEMENT + PARTICLE MOVEMENT] 4. BOTH LIQUID and GAS are PRESENT during the BOILING PROCESS [PRESENCE OF LIQUID + GAS] 5. The TEMPERATURE remains CONSTANT throughout the BOILING PROCESS until ALL LIQUID has BOILED [CONCLUSION] AS GAS HEATS UP: 1. After ALL LIQUID has boiled off, THERMAL ENERGY is converted to KINETIC ENERGY of the particles. [ENERGY CONVERSION] 2. The particles can move QUICKLY and RANDOMLY, in ANY DIRECTION [PARTICLE MOVEMENT] 3. The TEMPERATURE of the GAS rises BEYOND (BOILING POINT) [CONCLUSION] Conversion by Heat Loss: AS GAS COOLS: 1. KINETIC ENERGY of the particles is converted to THERMAL ENERGY, which is TRANSFERRED to the SURROUNDINGS [ENERGY CONVERSION] 2. With LESS KINETIC ENERGY, the particles SLOW DOWN [PARTICLE MOVEMENT SPEED] 3. The TEMPERATURE of the gas reduces to (BOILING POINT) [CONCLUSION] AS GAS CONDENSES: 1. The particles LOSE ENERGY to the SURROUNDINGS and the TEMPERATURE is at CONDENSATION POINT [ENERGY CONVERSION + CONDENSATION POINT] 2. The particles with LESS ENERGY are DRAWN CLOSER TOGETHER by the INTERMOLECULAR FORCES OF ATTRACTION between them [STRONGER IFOA] 3. The arrangement of the particles become LESS DISORDERLY [PARTICLE ARRANGEMENT] 4. Both GAS AND LIQUID are PRESENT during the CONDENSATION PROCESS [PRESENCE OF GAS AND LIQUID] 5. The TEMPERATURE remains CONSTANT throughout the CONDENSATION PROCESS until ALL GAS has CONDENSED [CONCLUSION] AS LIQUID COOLS: 1. After ALL GAS has CONDENSED, KINETIC ENERGY of the particles is converted to THERMAL ENERGY which is transferred to the SURROUNDINGS [ENERGY CONVERSION] 2. With LESS KINETIC ENERGY, the particles SLOW DOWN and SLIDE PAST each other FREELY throughout the liquid [PARTICLE MOVEMENT + SPEED] 3. The TEMPERATURE of the LIQUID reduces to (FREEZING POINT) [CONCLUSION] AS LIQUID FREEZES: 1. The particles LOSE ENERGY to the SURROUNDINGS and the TEMPERATURE is at FREEZING POINT [ENERGY CONVERSION] 2. The particles with LESS ENERGY are DRAWN CLOSER TOGETHER by the INTERMOLECULAR FORCES OF ATTRACTION between them [STRONGER IFOA] 3. The PARTICLE ARRANGEMENT becomes more ORDERLY [PARTICLE ARRANGEMENT] 4. Both SOLID AND LIQUID are present during the FREEZING PROCESS [PRESENCE OF SOLID AND LIQUID] 5. The TEMPERATURE remains CONSTANT throughout the FREEZING PROCESS until ALL the LIQUID has SOLIDIFIED [CONCLUSION] AS SOLID COOLS: 1. After ALL LIQUID has SOLIDIFIED, KINETIC ENERGY of the particles is converted to THERMAL ENERGY which is transferred to the SURROUNDINGS [ENERGY CONVERSION] 2. Particles can VIBRATE AND ROTATE about their FIXED positions [PARTICLE MOVEMENT] 3. The TEMPERATURE of the SOLID reduces below (FREEZING POINT) [CONCLUSION] Why does TEMPERATURE remain CONSTANT during change of state? During a change of state (e.g., melting, boiling, freezing), the temperature of a substance remains constant because the energy absorbed/released is used to break or form the intermolecular bonds between particles rather than increasing the kinetic energy of the particles, which would raise the temperature FYI: Temperature is a measure of the average kinetic energy of the particles in a substance Processes of Heat Changes: Physical State Solid Liquid Gas Changes into Solid Freezing Deposition Changes into Melting Condensation Liquid Changes into Gas Sublimation Evaporation Diffusion: NET movement of particles from regions of HIGHER concentration to LOWER concentration. This process occurs in mostly liquids and gases. Factors affecting Diffusion Rate: Temperature Particles with MORE KINETIC ENERGY at HIGHER TEMPERATURES move MORE QUICKLY. Particles in liquids and gases hence DIFFUSE QUICKER than solids. Stirring/Mixing Stirring and mixing INCREASES KINETIC ENERGY of particles, allowing them to MOVE QUICKLY and SPREAD out to areas of LOWER CONCENTRATION Particle Mass Particles of GREATER MASS require MORE KINETIC ENERGY to move at a given speed. Hence, heavier particles DIFFUSE MORE SLOWLY than lighter particles at any temperature (determined by calculating Mr). Examples of Diffusion (Possible Applications): Tea 1. Tea bag placed into water 2. Particles of tea diffuse out of bag into water 3. Particles of tea move away from bag towards regions in water having lower concentrations of tea 4. Eventually, concentration of tea is uniform throughout Perfume 1. Liquid perfume sprayed out of bottle 2. More VOLATILE substances VAPORISE almost immediately 3. Other substances fall as liquids onto surfaces of skin 4. Perfume vapours DIFFUSE AWAY from bottle, through air, eventually into our nose 5. At HIGHER temperatures, vapours diffuse FASTER, allowing us to detect fragrance MORE QUICKLY Atomic Structure Atom: The SMALLEST particle having CHEMICAL characteristics of an ELEMENT The Three Main Sub-atomic Particles: 1. Protons carry a POSITIVE (+) charge. 2. Neutrons carry NO NET electrical charge, found in the nucleus between protons. 3. Electrons carry a NEGATIVE (-) electrical charge and are extremely tiny particles that move very QUICKLY. Parts of an Atom Proton Neutron Electron Relative Mass 1 1 1/1840 Relative Charge +1 0 -1 Traits in a typical Equal to number - Equal to number atom of electrons of protons Location in Atom In the nucleus In the nucleus Electron Shells *Proton/Atomic number: Number of PROTONS in an atom’s nucleus *Nucleon/Mass number: Total number of PROTONS + NEUTRONS in the nucleus Nuclide Notation: Sub-atomic Particles in Ions: Ion: The particle formed when an atom or group of atoms GAINS/LOSES electron(s), but the number of protons and neutrons REMAIN THE SAME Ions: For example, this is a Calcium ion: It has 20 protons & 20 neutrons It has 18 electrons as it LOST 2 ELECTRONS to form a Calcium ion Isotopes: Atoms of the SAME element with SAME proton/atomic number but DIFFERENT nucleon/mass numbers due to DIFFERENT number of neutrons. Isotopes undergo the SAME KIND of chemical reactions but MAY have DIFFERENT PHYSICAL PROPERTIES like density & melting/boiling points. How Isotopes are represented: Element-X “Element” is the element name and “X” is the nucleon/mass number of the isotope E.g. chlorine-35 & chlorine-37, hydrogen-1 & hydrogen-2 How Sub-atomic Particles are Distributed in an Atom: Electron Shells and Energy Levels Electrons CLOSEST to nucleus -> lowest energy (innermost shell) Electrons FARTHEST away from nucleus -> highest energy (valence shell) Electron Shells and Electrons First/Innermost electron shell holds max. 2 electrons Second & Third electron shell holds max. 8 electrons each Fourth electron shell holds max. 2 electrons (by syllabus) If valence shell is not fully filled, the atom is REACTIVE and WILL participate in CHEMICAL REACTIONS E.g. Helium/Argon atoms have full valence shells, hence they are unreactive Electronic Configuration Format: x, y, z, … Where x, y and z are the number of electrons in each shell X being the number of electrons in the 1st shell Y being the number of electrons in the 2nd shell and so on Chemical Bonding Noble Gases and Their Electronic Configurations: Noble gases like helium and argon are elements found in the rightmost group (Group 18) of the Periodic Table. They… Have full valence electron shells Have noble gas electronic configuration (2, 8… 2, 8, 8… etc) Chemically unreactive Are monoatomic, existing as single atoms Atoms without Full Valence shells: The valence shells of the atom may undergo changes that allow them to become more STABLE They do so by chemically combining with other atoms to form bonds Atoms form bonds in 3 ways, Loss of electrons (Ionic/Metallic bonding) Gain of electrons (Ionic bonding) Sharing of electrons (Covalent bonding) Ionic Bonding: Ion Formation: Loss or Gain of Electrons Positive Ions/Cations Formed when an atom LOSES ONE or MORE electrons Positively charged as Protons > Electrons in number Have net positive charges, have noble gas electronic configuration E.g When a sodium atom loses an electron, it forms a sodium cation E.g When a magnesium atom loses electrons, it forms a magnesium cation Negative Ions/Anions Formed when an atom GAINS ONE or MORE electrons Negatively charged as Electrons > Protons in number Have net negative charge, have noble gas electronic configuration E.g Formation of a chloride ion The Ionic Bond and Compounds Ionic Bond: MUTUAL electrostatic attraction between ions of opposite charges Ions of opposite charges join by IONIC BONDING. There is MUTUAL ELECTROSTATIC FORCES OF ATTRACTION between the ions of opposite charges, making them move towards each other & remain in position E.G Formation of sodium chloride Ionic Compound: A neutral substance that consists of ions of opposite charges held together by electrostatic forces of attraction between positive and negative ions Ionic Compounds Have NO NET charge - the total positive charge from the cations must be EQUAL to the total negative charge from the anions Are formed between metals and non-metals Dot-and-Cross Diagrams of Ionic Compounds: RULES: 1. Determine electronic configuration of the elements 2. Determine how many electrons the metal atom needs to LOSE 3. Determine how many electrons the non-metal atom should GAIN 4. Determine the RATIO of metal to non-metal atoms 5. Electrons should NOT be touching each other 6. Ions should look SYMMETRICAL 7. FULL STRUCTURE WITH INNER ELECTRON SHELLS MUST BE DRAW UNLESS STATED OTHERWISE 8. After drawing, ensure sum of positive charges = sum of negative charges Examples: (valence electrons shown only) Sodium and Fluorine CHARGES are seen Clear DISTINCTION between Na electrons and F electrons with the use of dots and crosses Include a LEGEND stating which electron comes from which atom Ionic Structure: Positive and negative ions exert ELECTROSTATIC forces all around themselves; a positive ion that is already attached to a negative ion will still electrostatically attract other negative ions & negative ions electrostatically attract other positive ions Giant Ionic Crystal Lattice: Three-dimensional structure of alternating positive and negative ions Covalent Bonding: A covalent bond is formed by the sharing of a pair of valence electrons between two NON-METAL atoms, the atoms contribute one electron each The resulting covalent bond is a strong force of attraction between the NUCLEUS of an atom and the PAIR OF SHARED VALENCE ELECTRONS Molecules are formed when atoms are COVALENTLY BONDED to achieve a noble gas configuration of a FULLY FILLED valence electron shell Covalent Compound Structure: Simple Molecules: Have a SIMPLE MOLECULAR STRUCTURE Exist as SMALL DISCRETE MOLECULES Simple Molecular Structure WITHIN the molecules, ATOMS are held together by STRONG INTRAMOLECULAR COVALENT BONDS BETWEEN the molecules, the MOLECULES are held by WEAK INTERMOLECULAR FORCES OF ATTRACTION Macromolecules: Have a GIANT COVALENT STRUCTURE Consist of LARGE NETWORK OF ATOMS held together by STRONG COVALENT BONDS Diamond An ALLOTROPE/form of Carbon Each carbon atom FORMS STRONG COVALENT BONDS with FOUR other carbon atoms Graphite An ALLOTROPE/form of Carbon Carbon atoms are arranged in LAYERS Each carbon atom arranged in HEXAGONAL RINGS joining THREE other carbon atoms Each carbon atom is joined to other carbon atoms by STRONG INTRAMOLECULAR COVALENT BONDS BUT the INTERMOLECULAR FORCES OF ATTRACTION between LAYERS of carbon atoms are WEAK (graphite isn’t coal but it looks like it, so…) Silicon Dioxide, SiO2 Sand and Quartz are examples of substances made up of Silicon Dioxide Each SILICON ATOM is joined by STRONG COVALENT BONDS to FOUR OXYGEN ATOMS Each OXYGEN ATOM is joined by STRONG COVALENT BONDS to TWO SILICON ATOMS You can remember them with these acronyms! SIFOOX (Sea-fox): SIlicon atom -> FOur OXygen atoms OXYTWOSI (Oxy-two-see): OXYgen atom -> TWO SIlicon atoms Metallic Bonding: Metallic Bonding: Mutual ELECTROSTATIC ATTRACTION between POSITIVELY CHARGED METAL IONS and the delocalised SEA OF MOBILE ELECTRONS Giant Metallic Lattice Structure: Described as a LATTICE of POSITIVE METAL IONS surrounded by a SEA OF DELOCALISED ELECTRONS Metal atoms LOSE their valence electrons and become POSITIVELY CHARGED metal ions These valence electrons are then DELOCALISED, moving FREELY between the POSITIVELY CHARGED METAL IONS Structure and Properties of Materials Differences between Elements, Compounds & Mixtures: Elements (Pure) Compounds (Pure) Mixtures (Impure) Made up of… Only 1 element 2 or more elements 2 or more elements or compounds Combination Chemical Physical Method Formation Natural Chemical Reactions Physical Mixing Occurrence (usually) Ratio of Fixed composition by Variable composition Constituents mass by mass Properties Different from its Similar to its constituent elements constituent substances Melting & Fixed Fixed Over a range of Boiling Points temperatures Separation Cannot be Broken down into its Broken down into its Technique broken down constituents by constituents by into simpler chemical methods physical methods substances (e.g. electrolysis) (e.g. filtration) Examples Beryllium Magnesium Sulfate Air Manganese Iron(II) Hydroxide Hydrochloric Acid Silver Water Salt Solution Separation High amounts Little to no amount Energy (usually) (usually) Physical Properties of Ionic Compounds: Physical Property Detailed Explanation M.P. & B.P. HIGH melting and boiling points MOSTLY solids at room temperature Structure: GIANT IONIC CRYSTAL LATTICE Energy: A LOT of THERMAL ENERGY is required to overcome the forces of attraction Attraction: VERY STRONG ELECTROSTATIC forces of attraction between OPPOSITELY CHARGED ions Hence, they have high melting & boiling points Applications: Commonly used as refractories (heat-resistant materials) to LINE the INNER WALLS of furnaces Electrical Conduct electricity in AQUEOUS & MOLTEN state Conductivity Cannot conduct electricity in SOLID STATE Reason for Electrical Conductivity (Molten/Aqueous State): STRONG ELECTROSTATIC forces of attraction between OPPOSITELY CHARGED ions are OVERCAME Hence, MOBILE ions can carry an ELECTRIC CURRENT when POTENTIAL DIFFERENCE is applied Reason for Electrical Conductivity (Solid State): Ions are IMMOBILISED in GIANT IONIC CRYSTAL LATTICE structure by STRONG ELECTROSTATIC forces of attraction between OPPOSITELY CHARGED ions Hence, there are NO MOBILE ions to carry an ELECTRIC CURRENT when POTENTIAL DIFFERENCE is applied Solubility Most are SOLUBLE in water INSOLUBLE in organic solvents (alcohol, acetone) Volatility LOW volatility (low tendency to evaporate) Hardness HARD but BRITTLE Reasons for Its Hardness & Brittle Property: STRONG ELECTROSTATIC forces of attraction between OPPOSITELY CHARGED ions make ionic compounds RESISTANT to DEFORMATION Under enough force, ions move AWAY from their fixed lattice positions & ions of equal charge APPROACH each other REPULSIVE forces between ions of equal charge become LARGER than ATTRACTIVE forces Causes giant ionic crystal lattice structure to SHATTER, hence making them hard but brittle Physical Properties of Simple Molecular Substances: Physical Property Detailed Explanation M.P. & B.P. Low M.P. and B.P. (< 200°C) MOSTLY gaseous at room temperature ↑(Mr of molecule), ↑(IMFOA), ↑(M.P. + B.P.) Structure: SIMPLE MOLECULAR STRUCTURE Energy: SMALL AMOUNT OF THERMAL ENERGY required to overcome the Attraction: WEAK INTERMOLECULAR forces of attraction between MOLECULES Hence, they have low melting & boiling points Electrical Do NOT conduct electricity in ANY state Conductivity No MOBILE IONS/ELECTRONS to carry an ELECTRIC CURRENT when POTENTIAL DIFFERENCE is applied Applications: Covalent molecules; oils are used as electrical insulators in TRANSFORMERS and CIRCUIT BREAKERS Solubility SOLUBLE in organic solvents Most are INSOLUBLE in water Volatility Some have HIGH volatility Applications: Many perfumes & flavourings contain small amounts of volatile simple covalent molecules which diffuses to produce a smell Physical Properties of Giant Covalent Substances: Allotrope: Different forms of the same element different in their structural arrangement of atoms Physical Property Detailed Explanation M.P. & B.P. High M.P. and B.P. Solids at room temperature have an EXTENSIVE NETWORK of STRONG INTRAMOLECULAR covalent bonds, fixing atoms in fixed orderly arrangements Reasons for high M.P. and B.P: Structure: GIANT COVALENT STRUCTURE Energy: A LOT of THERMAL ENERGY is required to overcome the Attraction: EXTENSIVE NETWORK of STRONG intramolecular covalent bonds Hence, they have high M.P. and B.P. Electrical Do NOT conduct electricity (except Graphite) Conductivity ALL valence electrons are used to form INTRAMOLECULAR covalent bonds Hence, NO MOBILE ions/electrons can carry an ELECTRIC CURRENT when a POTENTIAL DIFFERENCE is applied For Graphite, Each carbon atom has ONE valence electron NOT used to form the intramolecular covalent bond These electrons are MOBILE and move between the LAYERS to carry an ELECTRIC CURRENT when a POTENTIAL DIFFERENCE is applied Solubility Most are INSOLUBLE in water & organic solvents Hardness VERY HARD (Except Graphite) Reasons for its hardness: Made up of only STRONG INTRAMOLECULAR covalent bonds Large amounts of energy is needed to break the GIANT COVALENT structure For Graphite, SOFT & SLIPPERY Made up of WEAK INTERMOLECULAR forces of attraction between carbon atom layers LOW amounts of energy needed to overcome them Gives it its ability to SLIDE over each other EASILY Applications: Diamond is HARD, used to COAT drill bits & tools Graphite is SOFT & SLIPPERY, used in pencil lead for WRITING & DRAWING as its layers are SHEARED OFF when pressed on paper, leaving a black mark Macromolecules: Polymers are macromolecules: 1. Natural polymers (Silk, Wool, Starch, Rubber) 2. Man-made polymers (Polyester, Polystyrene, Nylon, Plastic) Physical Property Detailed Explanation M.P. and B.P. MOST are solids at room temperature due to large size of polymers However, polymers may be formed by molecules of a RANGE OF SIZES Hence, they don’t have a fixed M.P. and B.P. Solubility MOST are insoluble in WATER Soluble in ORGANIC SOLVENTS Electrical Cannot conduct electricity due to absence of Conductivity mobile ions or electrons Hardness Molecular vibrations (high amount of kinetic energy) overcome the weak intermolecular forces of attraction between molecules. Hence, this causes softening over a wide range of temperatures Structure and Physical Properties of Metals and Alloys: Alloy: Mixture of a metal with ONE or MORE other elements Alloys have an IRREGULAR LATTICE ARRANGEMENT, leading to different properties from PURE metals which have a REGULAR LATTICE ARRANGEMENT Physical Property Detailed Explanation Malleability and Pure Metals: Ductility REGULAR lattice structure Layers of atoms SLIDE over one another EASILY when enough force is applied Easily bent/hammered into thin sheets Can be pulled into a wire without breaking Alloys: IRREGULAR lattice structure LARGER force needed to make the layers of atoms SLIDE over each other LESS malleable & ductile HARDER & stronger than pure metals M.P. and B.P. Pure Metals: HIGH M.P. and B.P. as atoms are held together in GIANT METALLIC LATTICES by STRONG metallic bonds Alloys: High M.P. and B.P. but melt and boil over a RANGE of temperatures as they are mixtures Thermal Pure Metals and Alloys: Conductivity GOOD conductors of thermal energy Sea of DELOCALISED electrons efficiently transfers thermal energy throughout the GIANT METALLIC LATTICE by conduction due to contacting of atoms Electrical Pure Metals and Alloys: Conductivity GOOD conductors of electricity Sea of DELOCALISED electrons carry ELECTRIC CURRENT when a POTENTIAL DIFFERENCE is applied Chemical Formulae and Equations Naming Ionic Compounds: 1. Positive ion comes BEFORE negative ion 2. The SECOND ion is usually a non-metal ion, ending with the name “-ide” [E.g reaction between oxygen with calcium metal -> Calcium (+) Oxide (-)] 3. No need for prefixes Naming Covalent Compounds: 1. Element with a smaller group number comes first 2. The SECOND element ends in “ide” 3. Need prefixes Compound name depends on: 1. TYPE of elements present 2. NUMBER of atoms of each element present No. of Atoms Prefix No. of Atoms Prefix 1 Mono- 4 Tetra- 2 Di- 5 Penta- 3 Tri- 6 Hexa- Writing Chemical Formulae: The VALENCY of an atom is the number of electrons involved, transferred in or given away, during bonding Valency = l Charge of Ion l (iykyk) Group 1 2 3-12 13 14 15 16 17 18 Valency 1 2 variable 3 4 3 2 1 0 Charge + +2 variable +3 +/- 4 -3 -2 - 0 TRANSITION METALS between Group 2 and 13 will have its valency indicated in the name of the compound, derived from its OXIDATION STATE. (E.G Chromium (IV) oxide, Valency = 4) These are the compounds you MUST memorise: Ion Formula Valency Charge Silver Ag 1 + Zinc Zn 2 +2 Hydroxide OH 1 - Nitrate NO3 1 - Carbonate CO3 2 -2 Sulfate SO4 2 -2 Ammonium NH4 1 + Understanding Chemical Equations: Format: Reactants -> Products Only write state symbols if the question requires it (s, l, g, aq) Forming Chemical Compounds: Balancing Chemical Equations: When forming chemical equations, BALANCING it is necessary There must be EQUAL numbers of each element before & after reaction Barium hydroxide + Ammonium chloride -> Barium chloride + Ammonia + Water Ba(OH)2 + NH4Cl -> BaCl2 + NH3 + H20 (Unbalanced) Ba(OH)2 + 2NH4Cl -> BaCl2 + 2NH3 + 2H20 (Balanced) Writing Ionic Equations: Must be BALANCED & state symbols MUST be present For AQUEOUS SUBSTANCES ONLY, re-write them as their CONSTITUENT ions Cancel out SPECTATOR IONS (‘Common terms’, for the math kiddos) E.G Sodium Carbonate + Nitric Acid -> Sodium Nitrate + Water + Carbon Dioxide A compound is AQUEOUS if it is SOLUBLE: Mole Concept and Stoichiometry Relative Masses: Relative Atomic Mass: Defined as the AVERAGE MASS of an atom of an element compared to 1/12 the mass of a Carbon-12 atom There are NO units for Ar Ar = Mass/Nucleon Number Relative Molecular Mass: Defined as the AVERAGE MASS of ONE molecule of that substance compared to 1/12 the mass of a Carbon-12 atom Mr = Sum of Ar of each elements in the substance Moles: What is the Mole?: 1 mole = (6.02 * 10^23) particles, known as Avogadro's constant Unit: mol Number of moles = Mass / Molar Mass Molar Mass: Refers to mass of one mole of atoms in the element Molar mass unit -> g/mol Mr = Molar mass [E.G Molar mass of O2 = 2(16) = 32 (Mr) = 32g/mol (MOLAR MASS)] Moles in Gas Calculations: Number of moles of gas = [Volume of gas (dm^3)] / 24dm^3 Volume of gas is proportional to number of moles Hence, Ratio of Moles of Gases = Ratio of Volumes of Gases Concentration of Solutions: No. Moles of Solution = Conc. of Solution * Volume of Solution Conc. of Solution Unit: mol/dm^3 or g/dm^3 Volume of Solution Unit: dm^3 Empirical & Molecular Formula: Empirical Formula SIMPLEST formula of a compound which shows the whole-number ratio of atoms of each element present in the substance E.G A sample of an oxide of copper contains 8g copper with 1g oxygen Molecular Formula ACTUAL number of atoms of each element in one molecule of the compound Relative Molecular Mass = n * Relative Mass of Empirical Formula E.G Propene has the EMPIRICAL FORMULA CH2 and its RELATIVE MOLECULAR MASS is 42. Find the MOLECULAR FORMULA (CH2)n = 42 [12 + 2(1)]n = 42 14n = 42 n=3 Stoichiometry: General rules: MAINLY focus on mole ratio Answers must ALWAYS be in decimals When doing working, leave it in 5 SF When presenting answer, leave it in 3 SF An example of an examination-styled Stoichiometry question is shown below. When Calcium Oxide is mixed with water, Calcium Hydroxide is formed. Calculate the maximum mass of Calcium Hydroxide which can be formed from 28 tonnes of lime (Calcium Oxide), given that 1 tonne is equivalent to 1000kg. Calcium Oxide + Water -> Calcium Hydroxide CaO + H2O -> Ca(OH)2 (Unbalanced) 2CaO + 2H2O -> 2Ca(OH)2 (Step 1: Form Balanced Equation) We have the mass of Calcium Oxide (Lime), hence we can make use of it to find the moles of Calcium Hydroxide by forming a MOLE RATIO 28 tonnes = 28 000 kg = 28 000 000 g No. of Moles in CaO = 28 000 000 / (40 + 16) No. of Moles in CaO = 500 000 mol Now we form the MOLE RATIO CaO : Ca(OH)2 2 : 2 1 : 1 500 000 : 500 000 With this, we can now find the mass of Calcium Hydroxide Mass of Ca(OH)2 = 500 000 * [ 40 + 2(16 + 1) ] Mass of Ca(OH)2 = 37 000 000g Mass of Ca(OH)2 = 37 tonnes Stoichiometry questions WILL be similar or slightly harder than this. Do practice questions on stoichiometry and it will be your free marks! Limiting & Excess Reactants: Many reactions carried out use an EXCESS amount of ONE reactant. The other reactant is completely used up (Limiting Reagent) Reactions STOP when the LIMITING REAGENT is completely used up The LIMITING REAGENT determines the amount of product formed Let’s take a look at a question that consist of the concept of Limiting Reagent: Hydrogen reacts with oxygen as shown in the equation below. How much gas will remain if 100cm^3 of hydrogen is reacted with 200cm^3 of oxygen at room temperature? 2H2 (g) + O2 (g) -> 2H2O (g) To start, we form a mole/gas ratio between the reactants and the product Remember, ratio of moles of gases = ratio of volumes of gases H2 : O2 : H2O 2 : 1 : 2 100 : 50 : 100 (possible) 400 : 200 : 400 (impossible, as we only have 100cm^3 of hydrogen) Hence, hydrogen is our limiting reagent Total gas at the end = (100 - 100) + (200 - 50) + 100 Total gas at the end = 250cm^3 Volumetric Analysis: Volumetric Analysis Uses a method called TITRATION Involves a solution of a KNOWN concentration (titrant) and a solution of an UNKNOWN concentration (analyte) Volume of titrant is used to determine the CONCENTRATION of the analyte, along with an indicator to show a VISIBLE colour change Colour changes happen when STOICHIOMETRIC AMOUNTS of both substances are added, indicating the COMPLETION of a reaction Percentage Calculations: Percentage Mass: Refers to the percentage mass of an element in a compound % mass = (Mr of element) / (Mr of compound) Percentage Purity: Often, a substance will not be pure due to the presence of impurities % Purity = (Actual mass of subst. / Given mass of subst.) * 100% Let’s take a look at an example of a question containing Percentage Purity 5.00g sample of copper powder was contaminated with copper (II) oxide. The copper (II) oxide in the sample was found to react with 0.02 mol of hydrochloric acid. Write the equation for the reaction and calculate the percentage of copper in the sample CuO + 2HCl -> CuCl2 + H2O It’s given that there is 0.02 mol of HCl, so let’s use it to find the ACTUAL mass of copper. We can do that by using a MOLE RATIO. HCl : CuO 2 : 1 0.02 : 0.01 Now let’s use the MOLE RATIO to find our ACTUAL mass of CuO! CuO mass = 0.01 * (64 + 16) CuO mass = 0.8g Now we can find our percentage of copper in the sample % purity = [(5.0 - 0.8) / 5.0] * 100% % purity = 84% (photo showing copper vs copper(II) oxide) Percentage Yield: Theoretical maximum mass is the mass of product calculated from the equation of the reaction Actual mass is the mass of the product obtained by conducting the experiment itself Actual mass is always lower than theoretical mass due to incomplete reactions, vaporisation of reactants/products and other experimental errors % Yield = (actual mass / theoretical max. mass) * 100% Let’s take a look at question involving percentage yield 28g copper(II) sulfate crystals were obtained from the reaction of 0.2 mol of copper(II) oxide. Calculate the percentage yield of copper(II) sulfate crystals) CuO + H2SO4 -> CuSO4 It’s given that there is 0.2 mol of copper(II) oxide, so let’s use that to form our mole ratio CuO : CuSO4 1:1 0.2 : 0.2 Mass of CuSO4 = 0.2 * [64+32+4(16)] Mass of CuSO4 = 32g % Yield = (28/32) * 100% % Yield = 87.5% Acids & Bases Acids: Properties of Acids: Substances producing hydrogen ions (H+ ions) when dissolved in water Exist as simple covalent compounds Have a SOUR taste H+ ions produced allow its aqueous solution to CONDUCT ELECTRICITY Turns blue litmus paper red React with REACTIVE metals, bases & metal carbonates Forms a resulting aqueous solution in water of pH values BELOW 7 Acids & Their Reactions: Acid + Reactive Metal -> Salt + Hydrogen E.g. Sulfuric Acid + Magnesium -> Magnesium Sulfate + Hydrogen Acid + Base -> Salt + Water E.g. Hydrochloric Acid + Sodium Hydroxide -> Sodium Chloride + Water Acid + Metal Carbonate -> Salt + Water + Carbon Dioxide E.g. Nitric Acid + Sodium Carbonate -> Sodium Nitrate + Water + Carbon Dioxide Solid acid salts remain as simple molecules when dissolved in organic solvents, producing NO H+ ions, resulting in the aqueous solution behaving unlike an acid. However, solid acid salts produce H+ ions upon dissolving in water. The presence of H+ ions allows the resulting aqueous solution to behave like an acid. Concentration vs Strength of Acids: Concentration: A measure of the amount of solute dissolved in a given volume of solvent Concentrated acid: Made by dissolving a LARGE amount of acid in a SMALL volume of water Dilute acid: Made by dissolving a SMALL mass/amount of acid/alkali in a LARGE volume of water Strength: The EXTENT of ionisation of acid molecules when dissolved in water Strong acid: Undergoes COMPLETE IONISATION in water to produce a LARGE concentration of hydrogen ions, H+ ions, in an aqueous solution No undissociated molecules remain [E.G dilute hydrochloric acid, sulfuric acid, nitric acid] Weak acid: Undergoes PARTIAL IONISATION in water to produce a SMALL CONCENTRATION of hydrogen ions, H+ ions, in an aqueous solution Most acid molecules do not ionise and remain in water [E.G ethanoic acid, carbonic acid, citric acid, lactic acid] Ionisation vs Dissociation: ‘Dissociation’ in water is used to describe the dissolving of alkalis in water while ‘Ionisation’ in water is used to describe the dissolving of acids in water Basicity of an Acid: Basicity refers to the MAXIMUM number of HYDROGEN ions produced by a molecule of acid when it is ionised in water E.G HCl (aq) -> H+ (aq) + Cl- (aq) = monobasic E.G H2SO4 (aq) -> 2H+ (aq) + SO42- (aq) = dibasic Bases/Alkalis: Base is ANY metal oxide/hydroxide, meaning that a base contains either oxide ions (O2-) or hydroxide (OH-) ions Can be soluble/insoluble in water, but most are insoluble Soluble bases are alkalis Properties of Alkalis: Substance that produces hydroxide ions (OH-) in water Strong alkalis DISSOCIATE COMPLETELY in water E.G sodium hydroxide, potassium hydroxide, Group 1 alkalis Weak alkalis DISSOCIATE PARTIALLY in water E.G aqueous ammonia, calcium hydroxide Bitter taste and soapy feel All alkalis dissolve in water to form solutions which conduct electricity due to the presence of mobile ions that can carry an electric current Have a pH value of 7 and turns moist red litmus paper blue Alkalis & Their Reactions Alkali + Acid -> Salt + Water E.G Hydrochloric Acid + Sodium Hydroxide -> Sodium Chloride + Water Alkali + Ammonium Salt -> Salt + Water + Ammonia E.G Sodium Hydroxide + Ammonium Chloride -> Sodium Chloride + Water + Ammonia In neutralisation reactions where acid & alkali react to form salt & water, hydrogen (H+) ions from acids react with hydroxide (OH-) ions from alkalis react to form water. *Hence the ionic equation for any neutralisation reaction is… H+ (aq) + OH- (aq) -> H2O (l) Comparing Relative Acidity and Alkalinity pH Scale Used to indicate whether solution is acidic/neutral/alkaline Ranges from pH 0 to pH 14 pH < 7: Acidic, solution has higher concentration of H+ ions pH 7: Neutral, solution has an equal number of H+ ions and OH- ions pH > 7: Alkaline, solution has higher concentration of OH- ions Indicators Substances that change colours in solutions of different pH levels Non-pH Indicator (E.g. Red & Blue Litmus Papers) ONLY INDICATES if solution is ACIDIC/ALKALINE Does not differentiate between WEAK and STRONG acids pH Indicator Indicates if solution is ACIDIC/NEUTRAL/ALKALINE Can determine how ACIDIC or how ALKALINE a solution is as different colours are observed at different pH levels Universal Indicator Mixture of different indicators Used to determine pH of solution Two types; UI Paper & UI Solution Natural Indicators These NATURALLY OCCURING indicators change colour to reflect the pH of the environments they are in E.G Lichen, Hydrangea, Turmeric, Red Cabbage, Brinjal, Blue Berries pH Meter/Datalogger with pH probe NOT an indicator More accurate in MEASURING pH than UI Measures pH to ONE or TWO decimal places More expensive than UI Determining Acidity & Alkalinity: A strength of an acid or an alkali can be determined with 3 ways: 1. Comparison of pH using pH indicator or pornhub meter 2. Measuring electrical conductivity 3. Determining RATE of reaction Measuring Electrical Conductivity Strong acids/alkalis have STRONGER electrical conductivity than weak acids/alkali of the SAME CONCENTRATION Strong acids/alkalis COMPLETELY IONISE in water, hence having HIGHER concentration of ions Weak acids/alkalis PARTIALLY IONISE in water, hence having LOWER concentration of ions Strong acids/alkalis have MORE mobile ions to carry an ELECTRICAL CURRENT when a POTENTIAL DIFFERENCE is applied Determining Rate of Reaction Strong acids reacts FASTER & MORE VIOLENTLY than weak acids of the SAME CONCENTRATION Produces a LARGER volume of H2/CO2 gas within the SAME TIME when reacting with REACTIVE metals/metal carbonates Controlling pH of Soils: It’s important to control soil pH as it affects growth and development of plants Most plants grow best when soil pH is around pH 6 to pH 7 Soil can become TOO ACIDIC from excessive use of chemical fertilisers/acid rain/microbial activities Excess acid in soil can be treated by LIMING Liming is the use of bases like CALCIUM HYDROXIDE (slaked lime) to neutralise the acidic soil by reacting with the acids, RAISING the pH level and promoting HEALTHY GROWTH However, adding excess Calcium Hydroxide will make the soil TOO ALKALINE & UNSUITABLE for plant growth Types of Oxides: Basic Oxides (Metal Oxides - solid at r.t.p.) Exhibits basic properties, reacting with acid to form salt & water Most basic oxides are insoluble in water However, Group 1 oxides dissolve readily in water, forming alkalis Amphoteric Oxides (Metal Oxides - solid at r.t.p.) Exhibits both acidic and basic properties Reacts with acid/bases to form salt & water Insoluble in water E.g. Zinc/Aluminium/Lead(II) (Z.A.P) oxide are amphoteric oxides Acidic Oxides (Non-metal Oxides - usually gases at r.t.p.) Exhibits acidic properties, reacting with bases to form salt & water Most acidic oxides dissolve readily in water, forming acids However, Silicon dioxide (SiO2) is insoluble in water E.g. Nitrogen Dioxide, Carbon Dioxide, Phosphorus Pentoxide Neutral Oxides (Non-metal Oxides - usually gases at r.t.p.) Exhibits neither acidic nor basic properties Insoluble in water E.g. Nitric oxide, Carbon monoxide & Water (NO CO H2O) are neutral oxides Salts Salts: An IONIC COMPOUND that consists of a CATION and an ANION Formed when a metallic/ammonium ion replaces one or more hydrogen ions of an acid during an acid reaction Salts produced from acid reactions MAY OR MAY NOT dissolve in water Preparation of Salts: There are 3 methods of preparing of any salt, depending on two main factors Solubility of salt in water Solubility of reagents in water General Rule: 1. If salt is insoluble: Precipitation 2. If salt is soluble, containing Grp 1/Ammonium ion: Acid + Excess Solid (MFC) 3. If salt is soluble w/o Grp 1/Ammonium ions: Titration *All 3 methods use reagents which contain the ions of the resulting salt* Precipitation: What it is and its General Rules: Used to prepare an INSOLUBLE salt by mixing 2 aqueous solutions Solution AB (aq) + Solution XY (aq) -> Insoluble Salt AY (s) + Solution BX (aq) Solution AB & Solution XY must be SOLUBLE as the ions must be able to MOVE and INTERACT with each other when reactants are mixed Standard Procedure of Precipitation: 1. Determine solutions AB and XY to be used for the reaction 2. Identify & mix Solution AB and Solution XY in excess in a BEAKER 3. Precipitate of Insoluble salt AY is formed 4. Filter mixture to obtain RESIDUE (Precipitate) 5. Rinse the precipitate with COLD DISTILLED water & dry by blotting between filter papers (Standard Procedure) Acid + Excess Solid (MFC): What it is and its General Rules: Used to prepare SOLUBLE salts that DO NOT contain Grp 1/Ammonium ions Acid AB (aq) + Substance XY (s/aq) -> Salt AY (s) + By-products (l/g) The reaction involves either of the following interactions: ○ Acid + Reactive Metal -> Salt + Hydrogen ○ Acid + Insoluble Base -> Salt + Water ○ Acid + Insoluble Carbonate -> Salt + Water + Carbon Dioxide In these reactions, solids reacting with acids must be in EXCESS & INSOLUBLE to COMPLETELY use up the acid & ensure that the soluble salt does not DISSOLVE into the excess solution which may affect filtration Standard Procedure of Acid + Excess Solid (Mix, Filter, Crystallisation): 1. Determine the solid in excess & acid to be used in the reaction 2. Add solid in EXCESS to HOT acid while stirring the mixture constantly until saturated (M) 3. Filter the mixture to remove excess solid, the filtrate contains the soluble salt (F1) 4. Heat the filtrate until it is SATURATED 5. Allow the filtrate to COOL, crystals of the soluble salt will form (C) 6. Filter to remove the crystals (F2) 7. Rinse the crystals with COLD DISTILLED water & dry by blotting between pieces of filter paper (Standard Procedure) NOTE: Reactive metals like sodium & potassium cannot be used as they react violently with acids, making the reaction dangerous to perform Titration: What it is and its General Rules: Used to prepare a SOLUBLE salt containing Group 1/Ammonium ions The reaction involves either of the following interactions: ○ Acid + Alkali -> Soluble Salt + Water ○ Acid + Soluble Metal Carbonate -> Soluble Salt + Water + Carbon Dioxide As both reactants are soluble, exact quantities must be used to avoid contamination of the final product The quantities can be determined by a suitable indicator, determining whether a reaction is complete as the reactants are usually colourless Standard Procedure of Titration: 1. Fill up Burette with an alkali/soluble metal carbonate solution 2. Pipette 25.0cm^3 of acid into a CONICAL flask 3. Add TWO drops of SUITABLE indicator to the acid 4. Add the alkali/soluble metal carbonate solution from the burette into the conical flask until the indicator JUST CHANGES colour PERMANENTLY 5. IDENTIFY EXACT volume of alkali/soluble metal carbonate solution used Preparation of Salt Solution: 6. Pipette 25.0cm^3 of acid into a NEW CONICAL flask 7. Add the SAME EXACT volume of alkali/soluble metal carbonate solution into the conical flask 8. Heat the solution until it is SATURATED 9. Allow the solution to cool, crystals of the soluble salt will form 10. Filter to obtain crystals, rinse the crystals with COLD DISTILLED water & dry by blotting between filter papers (Standard Procedure) Equivalence Point vs End Point Equivalence Point When pH 7 is reached (Neutralisation) End Point When the indicator changes colour as pH levels change A graph of pH shown in a pH meter: Ammonia The Haber Process: Ammonia is made industrially by the Haber Process, using a reversible reaction which can go both forwards and backwards simultaneously Nitrogen gas: Obtained from fractional distillation of liquid air Hydrogen gas: Obtained from the cracking of hydrocarbons or breaking down of crude oil fractions The chemical equation for the Haber Process is: N2 (nitrogen) + 3H2 (hydrogen) ⇌ 2NH3 (ammonia) Factors affecting the Haber Process: As the reaction is reversible, some Ammonia produced MAY BREAK DOWN & DECOMPOSE back into nitrogen & hydrogen Hence, CAREFULLY CONTROLLED pressure & temperature ensures the MAXIMUM yield of Ammonia at the MINIMAL cost HIGHER pressure leads to HIGHER yield of Ammonia & faster reaction but incurs a large cost (expensive equipment & large amount of electricity) LOWER temperatures REDUCES the DECOMPOSITION of Ammonia, leading to HIGHER yield but slows the reaction down -> SLOWER production Hence, pressure of 250 atm & temperature of 450°C is used The presence of a FINELY-DIVIDED iron catalyst is also used to further speed up the rate of reaction The Process of Ammonia Production: Obtaining Nitrogen & Hydrogen: 1. Nitrogen is obtained from fractional distillation of liquid air 2. Hydrogen is obtained from cracking of hydrocarbons in crude oil Production of Ammonia: 3. Nitrogen & Hydrogen are mixed in a 1:3 ratio by VOLUME 4. The resulting gas is COMPRESSED to a pressure of 250 atm 5. The compressed gas flows over an iron catalyst before heated to 450°C Collection & Storage of Ammonia: 6. Only about 15% of the leaving mixture is ammonia 7. A mixture of ammonia, hydrogen & nitrogen is then obtained & cooled 8. Ammonia gas condenses into a liquid, pumped into tanks & stored under pressure Recycling of Nitrogen & Hydrogen: 9. Unreacted nitrogen & hydrogen are then transferred back into the container to be recycled NOTE! At room temperature, nitrogen gas is unreactive, preventing a forward reaction to take place. Hence, a relatively high temperature & pressure are required to start a forward reaction between nitrogen & hydrogen Nitrogen & Hydrogen are mixed in a 1:3 ratio by volume, not by mole. This is due to the difference in relative atomic masses of nitrogen & hydrogen (14:3), resulting in a mole ratio unsuitable for production Oxidation & Reduction Gain & Loss of Oxygen: Oxidation: Gain of Oxygen The gain of Oxygen E.G Fe2O3 + 3CO -> 2Fe + 3CO2 [CO is oxidised, it gains 1 Oxygen atom to form CO2] Reduction: Loss of Oxygen The opposite of Oxidation; The loss of Oxygen E.G Fe2O3 + 3CO -> 2Fe + 3CO2 [Fe2O3 is reduced, it loses 3 Oxygen atoms to form Fe] Gain & Loss of Hydrogen: Oxidation: Loss of Hydrogen The loss of Hydrogen E.G CH4 (g) + 2O2 (g) -> CO2 (g) + 2H2O (g) [CH4 is oxidised, losing 4 Hydrogen & gaining 2 Oxygen atoms, forming CO2] Reduction: Gain of Hydrogen The gain of Hydrogen E.G 2H2 (g) + O2 (g) -> 2H2O (g) [Hydrogen gains Oxygen & is oxidised to water vapour] [Oxygen gains Hydrogen & is reduced to water vapour] Gain & Loss of Electrons: Oxidation: Loss of Electrons The loss of Electrons E.G Fe -> Fe(3+) + [3e-] [Every Iron atom loses 3 electrons] Reduction: Gain of Electrons The gain of Electrons E.G O2 + [4e-] -> 2O(2-) [Every Oxygen atom gains 2 electrons] Increase or Decrease in Oxidation State: Oxidation State: The charge an atom/element would have if existing as an ion in compound May take on a positive/negative integer, or zero Increase in oxidation state -> Loss of electrons -> Oxidation Decrease in oxidation state -> Gain of electrons -> Reduction Unchanged oxidation state -> No gain/loss of electrons -> Neither Oxidation nor Reduction occurred Calculating Oxidation States: Rule 0 (All elements): Atoms of the same element have no net charge (e.g. O2, C60) Oxidation state -> 0 Rule 1 (Group 1 and 2 metals): Group 1 element oxidation state -> +1 (loss of 1 electron) Group 2 element oxidation state -> +2 (loss of 2 electrons) Rule 2 (Special Non-Metals): Fluorine oxidation state -> -1 Hydrogen oxidation state -> usually +1 ○ When it forms compounds with metals, it’s oxidation state is -1 Oxygen oxidation state -> -2 ○ In peroxides, the oxidation state is -1 Rule 3 (Group 13 to 17 elements): Group 13 to 17 elements oxidation state -> often most common valency Metallic elements oxidation state -> almost always (+) Non-metallic elements & metalloids oxidation states -> may be (+/-) Rule 4 (Group 3 to 12 transition metals, lanthanides & actinoids): Group 3 to 12 elements oxidation state -> ALWAYS positive Roman numerals indicate the oxidation state present in the named species ○ Chromium(III) indicates that chromium has an oxidation state of +3 Identifying & Analysing Reactions: Redox Reactions: Reactions where both oxidation & reduction occurs Respiration C6H12O6 (aq) + 6O2 (g) -> 6CO2 (g) + 6H2O (l) [Carbon in C6H12O6 oxidises, increasing in oxidation state from 0 in C6H12O6 to +4 in CO2] [O2 reduces, decreasing in oxidation state from 0 in O2 to -2 in H2O] Photosynthesis 6CO2 (g) + 6H2O (l) -> C6H12O6 (aq) + 6O2 (g) [Carbon in CO2 reduces, decreasing in oxidation state from +4 in CO2 to 0 in C6H12O6] [Oxygen in H2O oxidises, increasing in oxidation state from -2 in H2O to 0 in O2] The Haber Process N2 (nitrogen) + 3H2 (hydrogen) ⇌ 2NH3 (ammonia) [N2 reduces, decreasing in oxidation state from 0 in N2 to -3 in NH3] [H2 oxidises, increasing in oxidation state from 0 in H2 to +1 in NH3] Non-redox Reactions: Reactions displaying no change in oxidation states Acid-Base Reactions (Neutralisation) E.g. Hydrochloric Acid + Potassium Hydroxide -> Potassium Chloride + Water HCl (aq) + KOH (aq) -> KCl (aq) + H2O (l) (Chemical Equation) H+ (aq) + OH- (aq) -> H2O (l) (Ionic Equation) *Hydrogen & Oxygen maintain an oxidation state of +1 & -2 respectively throughout the reaction, hence this is not a redox reaction. Precipitation of Salts E.g. Sodium Chloride (aq) + Silver Nitrate (aq) -> Sodium Nitrate (aq) + Silver Chloride (s) NaCl (aq) + AgNO3 (aq) -> NaNO3 (aq) + AgCl (s) (Chemical Equation) Ag+ (aq) + Cl- (aq) -> AgCl (s) (Ionic Equation) *Silver & Chlorine maintain an oxidation state of +1 & -1 respectively throughout the reaction, hence this is not a redox reaction. Oxidising & Reducing Agents: Oxidising Agents: Oxidises another substance while being REDUCED itself INCREASES the oxidation state of an element in another substance GAINS ELECTRONS from another substance E.g. Oxygen, Chlorine, Potassium Manganate(VII) Solution Reducing Agents: Reduces another substance while being OXIDISED itself DECREASES the oxidation state of an element in another substance LOSES ELECTRONS from another substance E.g. Hydrogen, Reactive Metals, Carbon, Potassium Iodide Solution Testing for Reducing & Oxidising Agents: Procedure of Testing for Oxidising Agents: Potassium Iodide Solution 1. Add a few drops of COLOURLESS aqueous potassium iodide (reducing agent) to a solution of an unknown substance 2. If the unknown solution is an OXIDISING AGENT, the colourless mixture turns YELLOW-BROWN or potentially DARKER results which show the precipitation of iodine as Iodide ions are being OXIDISED to Iodine 2I- (aq) -> I2 (aq) + 2e- [This is due to the increase in oxidation state of Iodine from -1 in 2I- to 0 in I2] Testing for Oxidising Agents: Iodine is not very soluble in water, precipitating as a black solid in LARGE amounts & pale yellow in SMALL amounts Spotting of potassium iodide on a strip of filter paper can also be used to test for oxidising agents in GASES. If the spot turns from colourless to yellow-brown when exposed to the gas, the gas is an oxidising agent STARCH may be further added to confirm the presence of iodine as it turns DARK BLUE upon the presence of iodine Procedure of Testing for Reducing Agents: Acidified Potassium Manganate(VII) 1. In this test, dark purple aqueous potassium manganate(VII) has been acidified with a little sulfuric acid, forming a purple solution 2. Add a few drops of the purple solution to the solution of an unknown substance 3. If the unknown solution is a reducing agent, the purple mixture turns colourless as manganate(VII) ions are REDUCED to manganese(II) ions 8H+ (aq) + MnO4- (aq) + [5e-] -> Mn(2+) (aq) + 4H2O (l) (acid) (purple sol.) (colourless sol.) [This is due to the decrease in oxidation state of Manganese from +7 in MnO4- to +2 in Mn(2+)] Testing for Reducing Agents: Spotting of acidified potassium manganate(VII) solution on a strip of filter paper can also be used to test for reducing agents in gases. If the spot decolourises when exposed to gas, the gas is a reducing agent NOTE: Reduction & oxidation occur simultaneously. If a reducing agent reacts with a substance, then the substance is an oxidising agent. Similarly, if an oxidising agent reacts with a substance, then the substance is a reducing agent Chemical Energetics Enthalpy Change: ΔH = Energy level of products - Energy level of reactants Exothermic Reactions: What is it? A reaction that releases energy from the system in the form of heat Exothermic Reaction Properties Temperature increases as energy is released from the reactants into the surrounding mixture Heat is lost from the reactants to the surroundings Energy of reactants > Energy of products ΔH < 0 (negative) as energy is lost to the surroundings [Energy Profile Diagram] [Effects of a Catalyst] Catalyst provides an alternative pathway with a LOWER activation energy for reaction to start. Hence, HIGHER PROPORTION of reactants will have sufficient energy to begin to react Activation energy is the minimum energy required by colliding reactant particles in order to react with each other Endothermic Reactions: What is it? A reaction that the system absorbs energy from its surrounding in the form of heat Endothermic Reaction Properties Temperature decreases as energy is absorbed by the reactants from the surrounding mixture Heat is gained from the surroundings by the reactants Energy of reactants < Energy of products ΔH > 0 (positive) as energy is absorbed from the surroundings [Energy Profile Diagram] [Effects of a Catalyst] Bond Breaking and Bond Formation: Some chemical reactions take in energy from the surroundings while others release energy into the surroundings ∆H = All energy absorbed to break bonds - All energy released to form bonds Bond Breaking Endothermic process Energy is absorbed from surroundings to break the bonds Total energy absorbed > Total energy released ∆H > 0 (positive) Bond Formation Exothermic process Energy is released into the surroundings form the bonds Total energy absorbed < Total energy released ∆H < 0 (negative) Bond Energy: Refers to the amount of energy absorbed or released to break or form one mole of a chemical bond Bond energies will be provided in the question Overall Energy Change = Sum of Reactant Energy - Sum of Product Energy The Periodic Table Periodic Table: A list of elements arranged in order of increasing proton/atomic numbers Divides elements into periods and groups ‘Groups’ refer to the vertical column of elements ‘Periods’ refer to the horizontal row of elements Proton Number & Electronic Configuration: Electronic Configuration Across a Period: Number of Electron Shells = Period Number Electronic Configuration Down a Group: Elements in the same Group have the same number of valence electrons Elements in the same group have similar chemical properties For elements in Groups 13 to 18, ○ Number of Valence Electrons = Group number - 10 Metallic & Non-Metallic Properties: Metallic Properties of Elements Across a Period: There is a DECREASE in metallic properties & an INCREASE in non-metallic properties across a period More energy is required to lose electrons (-) as there are more protons (+) in the nucleus An atom displays MORE metallic properties when it is MORE LIKELY to LOSE electrons than to GAIN electrons Metallic Properties of Elements Down a Group: There is an INCREASE in metallic properties & a DECREASE in non-metallic properties Size of the atom INCREASES down a group -> Valence electrons will be further away from the ATTRACTIVE FORCE of the nucleus. Hence, these valence electrons will be lost more easily Group 1 Elements - Alkali Metals: Physical Properties of Alkali Metals: Soft & can be cut easily (by a knife) Low densities, melting & boiling points ○ Going down the group, their densities GENERALLY increase (consists of some anomaly) ○ Going down the group, the melting points decrease Chemical Properties of Alkali Metals: Highly reactive Stored in oil to prevent reaction with air and water Achieves noble gas configuration by losing 1 valence electron Going down the group, ○ Size of atom increases -> Easier to lose valence electrons due to weaker attractive force of nucleus due to being further away from nucleus -> Reactivity increases Observations for Alkali Metals’ Reaction with Water: Lithium ○ Reacts quickly and floats on water Sodium ○ Reacts VIOLENTLY and DARTS around water surface, reaction MAY be explosive Potassium ○ Reacts VERY VIOLENTLY and is EXPLOSIVE Group 17 Elements - Halogens: Physical Properties of Halogens: Low melting and boiling points Coloured Going down the group, ○ Melting and boiling points increase ○ Colours become darker (Colour intensity increases) Chlorine ○ M.P. -> -101 ℃ ○ B.P. -> -34 ℃ ○ Appearance at R.T.P. -> Yellow-Green GAS Bromine ○ M.P. -> 7 ℃ ○ B.P. -> 59 ℃ ○ Appearance at R.T.P. -> Red-Brown LIQUID Iodine ○ M.P. -> 114 ℃ ○ B.P. -> 184 ℃ ○ Appearance at R.T.P. -> Purple-Black GAS Chemical Properties of Halogens: Reactive Non-Metals Gains 1 electron to achieve noble gas configuration Reacts with most metals to form salts; Halides Going down the group, ○ Size of atom increases -> More difficult for nucleus to attract one more electron due to weaker attractive force of nucleus due to being valence electrons being further away from nucleus -> Reactivity decreases Undergoes displacement reactions with Halide solutions ○ Displacement reaction is a reaction in which one element takes the place of another element in the compound MORE REACTIVE halogens will displace LESS REACTIVE halogens from its halide solution E.G Cl2 (aq) + 2KBr (aq) -> 2KCl (aq) + Br2 (aq) (light yellow) (colourless) (colourless) (red-brown) [Chlorine displaces bromine from a bromide solution] Group 18 Elements - Noble Gases: Properties of Noble Gases: Monoatomic Non-Metals Colourless gases at Room Temperature Low Melting and Boiling Points Insoluble in Water Unreactive ○ Noble gases have 8 valence electrons (Except Helium, 2 V.E) ○ As they have fully filled valence shells, they do not lose, gain or share electrons ○ Hence they rarely react to form compounds (This is the Noble M500, I think you can tell I’m bored) Properties of Transition Metals: High Melting Points and High Densities: Have higher melting points and densities than Group 1 and 2 Metals Variable Oxidation States: Just calculate the Oxidation State as learnt in Oxidation & Reduction (Remember the Sum of Oxidation States in a Compound is ALWAYS equal to 0) Coloured Compounds: Colours of Compounds of a Transition Metal are different at different oxidation states (TBC Memorise) Copper (II) -> Blue Iron (II) -> Pale Green Iron (III) -> Yellow / Brown Potassium Manganate (VII) -> Purple Catalysts: Substance that increases rate of chemical reaction and remains chemically unchanged at the end of the reaction ○ Iron -> Haber process for manufacture of ammonia The Reactivity Series Can be used to predict Behaviour of a metal from its position in the reactivity series Position of an unfamiliar metal in the reactivity series given a set of experimental results Potassium (Most Reactive) Sodium Calcium Magnesium Aluminium Carbon Zinc Iron Lead Hydrogen Copper Silver Gold (Least Reactive) The order of reactivity of metals can be determined by the reactions of metals with COLD WATER, STEAM or DILUTE HYDROCHLORIC ACID Reaction of Metals with Cold Water and Steam: Metal + Water -> Metal Hydroxide + Hydrogen Metals such as Potassium and Magnesium react with cold water More reactive metals react VIOLENTLY with cold water Metal + Steam -> Metal Oxide + Hydrogen Metals such as Zinc and Iron do not react with cold water but react with steam More reactive metals react VIOLENTLY with steam Observations for Reactions with COLD WATER: Potassium ○ Reacts VERY VIOLENTLY to form Potassium Hydroxide & Hydrogen gas ○ Enough heat is produced to cause Hydrogen gas to catch fire and EXPLODE Sodium ○ Reacts VIOLENTLY to form Sodium Hydroxide & Hydrogen Gas ○ Hydrogen gas may catch fire and EXPLODE Calcium ○ Reacts READILY to form Calcium Hydroxide & Hydrogen Gas Magnesium ○ Reacts VERY SLOWLY to form Magnesium Hydroxide & Hydrogen Gas ○ A test tube of Hydrogen gas is produced only after a few days (shows how slow the reaction is) Observations for Reactions with STEAM: Potassium, Sodium, & Calcium ○ Reacts EXPLOSIVELY Magnesium ○ Hot Magnesium reacts VIOLENTLY with steam to form Magnesium Oxide (white solid) and Hydrogen gas ○ A bright white glow is produced during the reaction Zinc ○ Hot Zinc reacts READILY with steam to produce Zinc Oxide and Hydrogen gas ○ Zinc Oxide is YELLOW when HOT and WHITE when COLD Iron ○ Red-hot Iron reacts SLOWLY with steam to form Iron Oxide and Hydrogen gas ○ Iron must be CONSTANTLY HEATED for the reaction to PROGRESS Reactions of Metals with Dilute Hydrochloric Acid: Metal + Dilute Hydrochloric Acid -> Metal Chloride + Hydrogen Indicates how reactive the metals are, more reactive metals react more VIOLENTLY Observations for Reactions with Dilute Hydrochloric Acid: Metals above Hydrogen in reactivity series react with Dilute Hydrochloric Acid to produce Hydrogen gas Potassium & Sodium ○ React EXPLOSIVELY Calcium ○ Reacts VIOLENTLY to give Calcium Chloride & Hydrogen Gas Magnesium ○ Reacts RAPIDLY to give Magnesium Chloride & Hydrogen Gas Zinc ○ Reacts MODERATELY FAST to give Zinc Chloride & Hydrogen Gas Iron ○ Reacts SLOWLY to give Iron (II) Chloride & Hydrogen Gas Reduction of Metal Oxides with Carbon: Metal Oxide + Carbon -> Metal + Carbon Dioxide Heat The lower the metal in the reactivity series from Carbon, the more readily the reduction of the metal oxide will occur Oxides of zinc require the HIGHEST TEMPERATURE for reduction Reduction of Metal Oxides with Hydrogen: Metal Oxide + Hydrogen -> Metal + Steam Hydrogen will reduce the oxide of metals from Iron and below Oxides of Iron will require the HIGHEST TEMPERATURE for reduction How Reactivity of Metals affect Tendency to form Positive Ions: A MORE REACTIVE metal has a GREATER tendency to form positive ions compared to a LESS REACTIVE metal ○ When Potassium reacts with water, it loses electrons READILY to form positive ions ○ When Magnesium reacts with water, it loses electrons LESS READILY ○ This is why Potassium reacts VIOLENTLY with water and Magnesium reacts VERY SLOWLY with water Displacement Reactions of Metals: A MORE REACTIVE metal can displace a LESS REACTIVE metal from its salt solution ○ When Iron filings are added to a solution of Copper (II) Sulfate ○ Copper metal is displaced out of the solution as a pink solid (when freshly formed) which then turns into a red-brown solid (after awhile) ○ The solution turns green Iron + Copper (II) Sulfate -> Iron (II) Sulfate + Copper 1. Iron has displaced Copper from Copper (II) Sulfate solution 2. A more reactive metal forms positive ions more readily 3. Since Iron is MORE REACTIVE than Copper, Iron atoms become Iron (II) ions and form Iron (II) Sulfate Considered a Redox Reaction ○ Magnesium reduces the Copper (II) ions to Copper ○ Magnesium itself is OXIDISED to the Magnesium ions Reaction between a Metal and the Oxide of Another Metal: A MORE REACTIVE metal can reduce the oxide of a LESS REACTIVE metal Displacement reaction The MORE REACTIVE metal displaces the LESS REACTIVE metal from its oxide Action of Heat on Metal Carbonates: The MORE REACTIVE a metal is, the MORE DIFFICULT it is to DECOMPOSE its carbonate by HEAT ○ MORE REACTIVE metals form carbonates that are MORE STABLE to heat than others Potassium Carbonate & Sodium Carbonate (higher up in reactivity series) ○ UNAFFECTED by HEAT Calcium, Magnesium, Zinc, Iron (II), Lead (II), Copper (II) Carbonates ○ Decompose into Metal Oxide and Carbon Dioxide on heating ○ Carbonates of metals below Sodium in the reactivity series decompose to form the Oxides of the metals and Carbon Dioxide Silver Carbonate ○ Decomposes into Silver and Carbon Dioxide on heating ○ The Silver Oxide produced is thermally unstable, hence it further decomposes to form Silver How Metals are Extracted from their Ores: Two main methods for extracting metals from their ores: 1. Reducing the metal compound to the metal using carbon 2. Electrolysis Reduction with Carbon Used for metals which are less reactive (Zinc and below) Reduction with Electrolysis Used for metals which are more reactive (Magnesium and above) The compounds are very difficult to break down and can only be extracted by Electrolysis Conditions for Rusting: The presence of both Oxygen (in air) and water are NECESSARY for rusting to occur Rusting is the slow oxidation of Iron to form Iron (III) Oxide Iron + Oxygen + Water -> Hydrated Iron (III) Oxide 4Fe (s) + 3O2 (G) + 2xH2O (l) -> 2Fe2O3,xH2O (s) Rust Prevention: Barrier Method: Prevents rusting by keeping Iron and Steel away from Oxygen (in air) and water ○ Painting ○ Oiling/Greasing ○ Coating with Plastic Sacrificial Protection: Protection of Iron and Steel against rusting using a MORE REACTIVE metal such as Zinc Zinc is MORE REACTIVE than Iron and can REACT IN PLACE of Iron ○ Galvanising or Coating with Zinc ○ Attaching a MORE REACTIVE metal such as Zinc or Magnesium to Iron or Steel Galvanising Involves dipping the Iron object into Molten Zinc A thin layer of Zinc is coated onto the Iron object It protects the Iron object by corroding in place of the Iron Attaching a More Reactive Metal Involves attaching blocks of MORE REACTIVE metal such as Zinc to the Iron object The MORE REACTIVE metal corrodes in place of the Iron object How does Sacrificial Protection work? Since Zinc is a MORE REACTIVE metal than Iron, it has HIGHER TENDENCY to form POSITIVE IONS Zinc atoms lose electrons in preference to Iron atoms and prevents Iron from forming Iron (III) Oxide More reactive metals will corrode instead of Iron As long as a more reactive metal is present, Iron will not rust Rusting as a Special Case of Corrosion: Iron and Steel: Rusting only occurs in Iron and Steel and it involves both water and oxygen (in air or water) The rust flakes off, exposing more metal to react Once rusting starts, it continues until the metal is COMPLETELY DESTROYED Other Metals: Also require water and oxygen (in air or water) to rust However, the Oxide layer forms a protective coating, preventing further reaction Good to Know: White Precipitates Reaction between ions: When two solutions containing different ions are mixed, they may react to form an insoluble solid, which appears as a white precipitate. Neutralisation reaction: When an acid and a base react, they can form salt and water. If the salt is insoluble, it can appear as a white precipitate. Double displacement reaction: When two solutions containing different ions are mixed, they may exchange partners, resulting in the formation of a white precipitate. Bases vs Alkali Anything insoluble in water does not change the pH level of the water In order for the pH level to change, the substance must dissolve in the water to produce hydroxide/hydrogen ions Insoluble bases automatically do not change pH levels of water as they cannot even dissolve inside it in the first place Ionic Compounds As long as the compound contains elements of oppositely charged ions, it automatically is an ionic compound

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