Metal Hydroxide Precipitates and Solubility

• The notes provided cover Year 9 Chemistry Test 2 Revision, including flame tests, precipitate tests, reactions, and balancing equations.

Flame Tests for Cations

• Flame tests identify metal ions by the color they emit when heated strongly.
• Method involves dissolving the ionic compound, soaking a wooden splint in the solution, and burning the splint in a Bunsen burner flame.
• The flame's color is observed to identify the metal cation.

Summary of Flame Test Colors:

• Lithium ions (Li+) produce a pink-red flame.
• Sodium ions (Na+) produce a yellow flame.
• Potassium ions (K+) produce a lilac flame.
• Calcium ions (Ca2+) produce an orange-red flame.
• Barium ions (Ba2+) produce an apple-green flame.
• Copper ions (Cu2+) produce a green/blue flame.
• Under blue cobalt glass, sodium appears the same, while potassium, calcium, and barium appear white, and copper appears cyan/white.

Silver Precipitate Test for Halide Anions

• This test identifies halide ions by the color of the silver halide precipitate formed.
• This method involves adding silver nitrate solution (AgNO3(aq)) to the suspected halide solution and observing the precipitate color.

Summary of Silver Halide Precipitate Colors:

• Chloride ions (Cl-) produce a white precipitate.
• Bromide ions (Br-) produce a cream precipitate.
• Iodide ions (I-) produce a yellow precipitate.

Using Ammonia Precipitates to Identify Metal Ions

• This method distinguishes metal ions based on the solubility of their ammonia precipitates.
• The method involves taking metal salt solutions, adding ammonia solution (NH3(aq)), shaking, recording the precipitate appearance.

Summary of Ammonia Precipitate Test Results:

• Zinc sulfate (ZnSO4) forms a milky white, cloudy, grainy precipitate that is soluble in excess ammonia.
• Copper sulfate (CuSO4) forms an opaque, blue, grainy precipitate that is soluble in excess ammonia.
• Calcium chloride (CaCl2) forms a transparent, clear, grainy precipitate that is soluble in excess ammonia.
• Iron(II) sulfate (FeSO4) forms a grey-green, fine precipitate that turns dark green and is insoluble in excess ammonia.
• Iron(III) chloride (FeCl3) forms an orange-brown, fine, grainy precipitate that is insoluble in excess ammonia.
• Aluminum sulfate (Al2(SO4)3) forms a translucent, cloudy precipitate that is insoluble in excess ammonia.
• Chromium(III) sulfate (Cr2(SO4)3) forms a dark green-blue, fine precipitate that is insoluble in excess ammonia.
• Aluminium and zinc are similar; zinc is if it's soluble, and aluminium if it's not in excess ammonia.

Using Hydroxide Precipitates to Identify Metal Ions

• Hydroxide precipitates can identify metal ions based on the precipitates' solubility.
• To identify, use sodium hydroxide solution in a metal salt solution.
• If precipitate is formed, then it must be the metal hydroxide
• The method involves adding sodium hydroxide solution (NaOH(aq)) to the metal salt solution, shaking, and adding a test tube of the same NaOH(aq) for solubility testing.

Summary of Hydroxide Precipitate Test Results:

• Zinc sulfate (ZnSO4) forms a white and quite gelatinous, cloudy precipitate [Zn(OH)2] that is soluble (Zn2+).
• Copper sulfate (CuSO4) forms a blue and gelatinous precipitate [Cu(OH)2] that is insoluble (Cu2+).
• Calcium chloride (CaCl2) forms a liquid, white, cloudy, grainy precipitate [Ca(OH)2] that is insoluble (Ca2+).
• Iron(II) sulfate (FeSO4) forms a grainy and dark green precipitate [Fe(OH)2] that is insoluble (Fe2+).
• Iron(III) chloride (FeCl3) forms a reddish-brown, slightly gelatinous and grainy precipitate [Fe(OH)3] that is insoluble (Fe3+).
• Aluminum sulfate (Al2(SO4)3) forms a partly cloudy, clear, gelatinous precipitate [Al(OH)3] that is soluble (Al3+).
• Chromium(III) sulfate (Cr2(SO4)3) forms an emerald, liquidy, grainy, gelatinous precipitate [Cr(OH)3] that is soluble (Cr3+).

Solubility of Hydroxides

• All hydroxides will form insoluble hydroxides.
• Zinc, aluminum, and chromium are soluble hydroxides.
• The iron(II) hydroxide precipitate will have red particles start to form on the side of the test tube after some time.
• Aluminium precipitates do not last but redissolve when an excess of NaOH(aq) is added.
• Sodium hydroxide (NaOH) and ammonia are used to test these ions to differentiate between compounds by adding either compounds in excess.
• Copper is soluble in ammonia excess while Cr3+ and Fe2+ are both green and aren't soluble in ammonia excess
• Cr3+ is in sodium hydroxide

Covalent Bonding

• Covalent bonding happens due to the attraction of a shared pair of electrons.
• Two positive nuclei of atoms occur when neither atom has particularly stronger attraction than the other.
• No transfer of electrons takes place.
• Electrons are shared.

Naming Covalent Compounds:

• Write the name of the first non-metal in the chemical formula.
• Write the other non-metal's name with an "-ide" at the end.
• Add number prefixes to indicate the number of atoms.

Drawing Covalent Diagrams

• Only include electrons in the outer orbit.

Examples of Covalent Bonding:

• Hydrogen Fluoride (HF): Hydrogen (1 electron) shares with Fluorine (7 electrons) to attain full outer shells, forming a single bond.
• Chlorine (Cl2): Two chlorine atoms (7 electrons each) share one electron to form a single bond (Cl-Cl).
• Oxygen (O2): Oxygen atoms (6 electrons each) share two electrons to form a double bond (O=O).
• Nitrogen (N2): Nitrogen atoms (5 electrons each) share three electrons to form a triple bond (N≡N).
• Carbon Dioxide (CO2): Carbon (4) shares two electrons with each Oxygen (6) to form the double bonds (O=C=O).
• Methane (CH4): Carbon (4) shares one electron with each Hydrogen (1) to form four single bonds.

Properties of Bonds

• The maximum number of bonds between two atoms is three.
• The maximum number of bonds an atom can form with other atoms is four.

Ionic vs. Covalent Bonds

Property Differences:

• Ionic compounds consist of metals and nonmetals, while covalent compounds consist of nonmetals only.
• Ionic bonding involves attraction between anions and cations, while covalent bonding involves sharing pairs of electrons.
• The representative unit for ionic compounds is a formula unit, while for covalent compounds, it's a molecule.
• The formula represents the ratio of ions in ionic compounds and the type/number of atoms in covalent compounds.
• Ionic compounds are solid at room temperature, while covalent compounds can be gas, liquid, or solid.
• Ionic compounds usually have high water solubility; covalent compounds have variable solubility.
• Ionic compounds typically have high melting and boiling temperatures, while covalent compounds generally have low ones.
• Ionic compounds are good electrical conductors when molten or in solution, while covalent compounds are poor conductors.
• Ionic compounds separate into ions when dissolved in water, while covalent compounds remain as molecules.

Tests for Molecules

Carbon Dioxide Test:

• Passing the gas through limewater (calcium hydroxide solution)
• Limewater reacts with carbon dioxide to produce a precipitate of white calcium carbonate.
• Limewater looks milky.

Oxygen Test:

• A glowing splint will relight in the a test tube filled with 100% oxygen.

Hydrogen Test:

• Hydrogen burns to give water when hydrogen is mixed with air.
• A lit wooden splint will create a 'pop' sound.

Ammonia Test:

• Ammonia is the only alkaline common gas.
• Use moist litmus or indicator paper turns red litmus paper blue, indicating an alkaline nature.

Water Test:

• Anhydrous cobalt chloride turns pinkish-white.
• Anhydrous copper sulfate turns blue.

Chlorine Test:

• Chlorine is a very powerful oxidizing agent and will react with and bleach many colored compounds.
• Blue litmus paper put on tile turns red after a drop of chlorine water is added.

Iodine Test:

• Iodine mixed with starch solution creates a blue-black complex.

Atoms, Moles, and Molecules:

Atom:

• The smallest, electrically neutral particle of an element involved in a chemical change.

Molecule:

• The smallest, electrically neutral particle of an element or compound existing on its own.

Ion:

• An atom or group of atoms carrying an electric charge.
• A single oxygen atom (O) cannot exist on its own but can as part of a molecule (O2).
• Few compounds exist as a single atoms
• An atom, therefore, is the same as a molecule

Relative Atomic Mass (Ar):

• The weighted average of the masses of an element's isotopes compared to carbon-12, with no units.

Relative Molecular Mass (Mr):

• The sum of the relative atomic masses of atoms in the chemical formula.
• For H₂O: (1x2) + 16 = 18.

Relative Formula Mass:

• Simplest ratio of the atoms in the compound, usually referring to ionic substances.

Mole:

• A unit to convert relative mass into a measurable quantity.
• Each mole contains 6.02 × 10^23 particles (Avogadro's number).
• Avogadros number is the number of atoms in 12g of Carbon-12.

Molar Masses:

• Calculate by multiplying the number of atoms by their mass and adding those.
• Molecules contain atoms that make up Xg mol^-1, with the g mol^-1 being = per atom.

Calculate mass:

• By multiplying a mole of atoms by Avogadros number

Rates of Reaction:

• Reaction rates are affected by three factors:
  • Temperature(Heating increases reaction rates).
  • Concentration (Increase reactant concentration to increase reaction rates).
  • Surface Area: (Crushing solid lumps into a powder increases reaction rates).

Collision Theory:

• Particles must collide with sufficient energy to break bonds
• There must be a collision between reactant particles and must have sufficient energy

Observed Changes:

• In order to follow a reaction, observe changes that take place

Examples Include:

• Formation or dissolving of a precipitate
• Formation or dissolving of a colored substance
• Formation of a mass

Comparisons:

• When comparing the rates of a particular reaction under different conditions determine the time required for a fixed amount of reaction to occur

Surface Area Calculations

• An increase in surface area results in a faster reaction rate

Following Reactions:

• Gas production, surface area, concentration, and temperature can be used to measure the reaction rate.

Balancing Chemical Equations:

• Four main rules to follow:
  • Verify all formulae are correct.
  • Address one element at a time.
  • Use big numbers, not small, for balancing.
  • Review and revise each element by repeating the initial rules if needed.
• If the equation is unbalanced, add number of coefficient to balance.

Effects of Catalysts:

• Catalysts increase the reaction rate without being used but decrease the activation energy required.
• The use of catalysts can also have varying effects on the rate if they cannot be used up.