Ionic Bonding: Formation and Properties

Ionic Bonding

• Ionic bonds form when a metal reacts with a non-metal, resulting in the creation of ions.
• An ionic bond is a strong electrostatic attraction between a positive ion (cation) and a negative ion (anion).
• Multiple ions combine to create a giant lattice structure held together by numerous strong ionic bonds.

Formation of Ionic Bonds

• Metal atoms transfer outer shell electrons to non-metal atoms.
• Metals become positively charged cations with more protons than electrons.
• Non-metals become negatively charged anions with more electrons than protons.
• Atoms achieve greater stability by attaining a full outer shell of electrons.
• Sodium (2,8,1) loses an electron becoming Na+(2,8).
• Chlorine (2,8,7) gains an electron becoming Cl- (2,8,8).
• Magnesium Mg(2,8,2) loses 2 electrons, bonding with Oxygen O(2,6) which gains 2, to form MgO.
• Two Sodium atoms Na(2,8,1) each loses 1 electron, bonding with Oxygen O(2,6) which gains 2, to form Na2O.

Physical Properties of Ionic Compounds

• Ionic compounds form a giant lattice with many strong ionic bonds, thus have high melting points as lots of thermal energy is needed to break the bonds.
• NaCl has a melting point of 801 °C.
• MgO has a melting point of 2852 °C.
• Ionic bond strength increases with the charge of the ions (MgO has stronger bonds because of Mg2+ and O2-).
• Solid ionic compounds do not conduct electricity because ions are held in place.
• Molten or aqueous solutions of ionic compounds conduct electricity because ions are free to move.

Formulae

• Chemical formulae list the types and numbers of atoms in a substance.
• Molecular formulae specify the number of each element's atoms in a molecule.
• Metals form positive ions (cations) by losing electrons, whereas non-metals gain electrons to form negative ions (anions).
• Sodium (Na+) is a Group 1 ion with a +1 charge.
• Beryllium (Be2+) is a Group 2 ion with a +2 charge.
• Aluminium (Al3+) is a Group 3 ion with a +3 charge.
• Fluoride (F-) is a Group 7 ion with a -1 charge.
• Oxide (O2-) is a Group 6 ion with a -2 charge.
• Nitride (N3-) is a Group 5 ion with a -3 charge.
• Hydrogen (H+), Ammonium (NH4+), and Copper(I) (Cu+) are other +1 ions to memorize.
• Copper(II) (Cu2+), Iron(II) (Fe2+), and Lead(II) (Pb2+) are other +2 ions to memorize.
• Iron(III) (Fe3+) is another +3 ion to memorize.
• Hydroxide (OH-), Nitrate (NO3-), Sulphate (SO42-), and Carbonate (CO32-) are other negative ions to memorize.
• "-ide" ions consist of a single element, while "-ate" ions involve an element bonded to oxygen.
• Roman numerals indicate the charge of a metal ion when it can vary.

Simple Molecules

• Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Bromine (Br2), and Iodine (I2) are diatomic elements.
• Water (H2O), Carbon Monoxide (CO), Carbon Dioxide (CO2), Ammonia (NH3), and Methane (CH4) are simple compounds.
• Hydrochloric acid (HCl), Nitric acid (HNO3), and Sulfuric acid (H2SO4) are common acids.
• (s) = solid, (g) = gas, (l) = liquid, (aq) = aqueous solution; are state symbols.

Formula of Ionic Compounds

• Ionic compounds are named with the positive ion first, followed by the negative ion.
• The total positive charge must equal the total negative charge in the formula.
• Sodium Chloride (NaCl) is formed when one Na+ ion combines with one Cl- ion.
• Magnesium Chloride (MgCl2) is formed when one Mg2+ ion combines with two Cl- ions.
• Aluminium Oxide (Al2O3) is formed when two Al3+ ions combine with three O2- ions.
• Ammonium Sulfate ((NH4)2SO4) is formed when two NH4+ ions combine with one SO42- ion; use brackets if more than one molecular ion is needed.
• Selected examples: Sodium sulfate (Na2SO4), Calcium oxide (CaO), Magnesium nitrate (Mg(NO3)2), Iron(III) sulfate (Fe2(SO4)3), Sodium nitrate (NaNO3), Calcium carbonate (CaCO3), Aluminium nitride (AlN), Potassium iodide (KI).

Writing Equations

• Law of Mass Balance: Mass is conserved in chemical reactions.
• The number of atoms for each element must be equal on both sides of the arrow.
• Writing Chemical Equations:
  • Write the word equation.
  • Write the chemical formulas underneath each name.
  • Balance the equation using coefficients.
  • Add state symbols when needed.
• Example: Magnesium + Oxygen -> Magnesium Oxide shown by, 2Mg(s) + O2(g) -> 2MgO(s)
• Example: Nitric Acid + Calcium Carbonate -> Calcium Nitrate + Carbon Dioxide + Water, shown by 2HNO3(aq) + CaCO3(s) -> Ca(NO3)2(aq) + CO2(g) + H2O(l)
• Example: Methane + Oxygen -> Carbon Dioxide + Water: CH4(g) + 2O2(g) -> CO2(g) + 2H2O(l)

Metal Reactivity and Extraction

• Metals higher in the reactivity series lose electrons more easily (more reactive).
• Potassium (K) is most reactive, followed by Sodium (Na), Lithium (Li), Calcium (Ca), Magnesium (Mg), Aluminium (Al), CARBON (C), Zinc (Zn), Iron (Fe), HYDROGEN (H), Copper (Cu), Silver (Ag), and Gold (Au).
• Metals lower in the series lose electrons less easily (less reactive).
• Oxidation is loss of electrons; Reduction is gain of electrons (OIL RIG).
• Oxidizing agents accept electrons, causing oxidation.
• Reducing agents give away electrons, causing reduction.
• A more reactive element displaces a less reactive one in a compound.
• Copper oxide + Zinc -> Zinc oxide + Copper; here Copper oxide is reduced, and Zinc is oxidised.
• Copper sulfate + Zinc -> Zinc sulfate + Copper; the copper ions are reduced, and the zinc is oxidized.

Reaction of Metals with Water and Acids

• Metals above hydrogen in the reactivity series react with water/acids via displacement of hydrogen.
• Metal + Water -> Metal Hydroxide + Hydrogen; e.g Calcium + Water -> Calcium Hydroxide + Hydrogen Ca(s) + 2H2O(l) -> Ca(OH)2(aq) + H2(g).
• Metal + Acid -> Salt + Hydrogen (MASH); e.g. Calcium + Hydrochloric Acid -> Calcium Chloride + Hydrogen, Ca(s) + 2HCl(l) -> CaCl2(aq) + H2(g) or Sodium + Sulfuric Acid -> Sodium Sulfate + Hydrogen 2Na(s) + H2SO4(l) -> Na2SO4(aq) + H2(g).
• Faster reactions indicate more reactive metals.

Metal Extraction

• Metal ore: rock with enough metal for economical extraction, often as compounds with oxygen/sulfur.
• Platinum, gold, and silver are often found as native metals.
• Metals below carbon extracted via carbon reduction. Includes zinc, iron, and copper ores.
• Iron(III) oxide + carbon -> iron + carbon dioxide shown by Fe2O3(s) + 3C(s) -> 2Fe + 3CO2(g)
• Metals above carbon require electrolysis; Aluminium extracted this way, Al3+ + 3e- -> Al

Uses of Metals

• Aluminium: High Strength to weight ratio, resists corrosion so used for aircraft, power lines etc
• Copper: Resists corrosion thus used for cooking pans, and is high electrical conductivity, anti-microbial for electrical wires, water pipes etc.
• Iron: High strength when alloyed, and used for different types of steel.

Corrosion

• Iron corrodes and forms rust (hydrated iron oxide) in the presence of oxygen and water, accelerated by salt.
• Rusting prevention:
  • Zinc Coating (Galvanising) as Zinc is more readily oxidised
  • Sacrificial protection such as using blocks of more reactive metal
  • Barrier coating with oil, grease, paint, plastic or rubber.

Alloys

• Alloy definition: A metal mixed with another element, often a metal.
• Pure Metals: Atoms in layers that slide easily.
• Alloy Properties: Atoms of different sizes disrupt layers so alloys are harder than pure metals.

Types of Steel

• Low Carbon Steel: Iron with up to 0.25% carbon to make it malleable.
• High Carbon Steel: Iron with 0.6-1.2% carbon to make it hard and wear-resistant for knives.
• Stainless Steel: Iron with chromium and nickel for corrosion resistance, and used in cutlery.

Covalent Bonds

• Covalent bonds involve sharing electrons between non-metals.
• They are the electrostatic attraction between shared negative electrons and positive nuclei.
• Molecules are particles with two or more chemically bonded atoms where covalent substances are simple molecules.

Dot and Cross Diagrams

• Atoms share electrons to achieve a full outer shell by bonding with a partner.
• Shared pair of electrons symbolised with a line in a displayed formula, such as seen in H–H in H2 as a single covalent bond.
• Double covalent bonds (two shared pairs) are shown with two lines such as in O=O in O2.
• Triple covalent bonds (three shared pairs) are shown with three lines such as in N≡N in N2.

Properties of Covalent Molecules

• Covalent bonds are strong while intermolecular forces of attraction are weak leading to, compounds having Low Melting and Boiling points as boiling needs little energy to break weak intermolecular forces.
• Covalent molecules do not conduct electricity as they lack mobile charges (delocalised electrons or ions).

Giant Covalent Structures

• Applies to carbon (graphite, diamond, graphene), silicon, and silicon dioxide
• Diamond: Carbon atoms form strong tetrahedral bonds, hence diamond has a high of Melting point.
• Graphite: Carbon atoms atoms bond to 3 others in layers. Hence graphite has high melting point, is a Lubricant, and an Electrical Conductor because of the weak forces of attraction layers.
• Silicon: Similar to diamond hence also a Hard Semiconductor
• Silicon dioxide: A diamond-like silicon lattice giving is a hard, high melting point
• Graphene: Single sheet of hexagonally arranged carbon atoms giving is high strength to weight ratio, and an electrical conductor

Giant Covalent Structure Explanations

• Diamond's hardness arises from its rigid tetrahedral structure, so significant energy is needed to break bonds.
• The high melting points, result from giant structures of strong covalent bonds that need a lot of thermal energy to break.
• Graphite conducts electricity as carbon atoms use three outer electrons for covalent bonds, thus leaving delocalised electrons to move when a voltage is applied.
• Graphite is a lubricant due to weak attractions between layers, allowing them to slide.
• Water has weak intermolecular forces which do not require lots of energy to break, making it easy to melt.

Metallic Bonding

• Metals have Giant Metallic Structure, and metal atoms easily lose outer electrons to become delocalised resulting in strong electrostatic attraction between positive ions and delocalised electrons.

Properties of Metals

• Good electrical conductors as delocalised electrons move to carry charge.
• Malleable because layers of metal ions easily slide.

Relative Atomic Mass

• Definition: a weighted average of the mass of an atom when compared to 1/12 the mass of a Carbon-12 atom.

Relative Formula Mass

• Definition: the sum of the relative atomic masses of all the atoms in its formula
• For H2O; Mr = 1 + 1 + 16 = 18

Mole

• Definition: amount of substance with 6.022 x 10^23 particles (Avogadro's number)
• The Mr of a compound in grams is one mole
• Mass/mole formulas
  • Moles(n) = Mass(m) / Mr
• Calculations from equations- work out the moles of the substance you know the mass of, and use ratio to work out the mass of what you require.

Empirical Formula

• Definition: smallest version of the ratio of elements in a compound where each element has a whole number
• Definition: actual number of each element in a molecule
• Calculating
  • List the elements
  • Write the mass number from the experiment.
  • Divide each mass by the Ar
  • Divide the ratio with the smallest
  • Write the formula

% Water of Crystallisation

• Hydrated salt can be dehydrated to give anhydrous salt + water
• CuSO45H2O(s) -> CuSO4 + 5H2O(g)
• The amount of moles of the water and CuSO4 are in proportion with the ratio to find x.

Moles of Gases

• At room temperature and pressure, 1 mole has a volume of 24dm^3
• n = volume / 24dm^3

Percentage Yield

• Equation: Actual / Theoretical x 100
• When the yield is lowered in a chemical reaction
  • Side reactions and impurities

Excess and limiting Reagents

• A Reagent is considered in excess if there is an amount leftover by the end of the reaction.
• Steps:
  • Na + Cl2 -> NaCl
  • If Cl2 > Na, this makes Cl2 the excess and Na, the limiting Reagent.
  • %Yield = Actual mass obtained/ calculated mass x 100

Periodic Table

• Vertical columns (groups), horizontal rows (periods).
• Elements in order of Atomic Number (protons) and the group number = the number of outer shell electrons.
• Nobel Gases full outer shells are very unreactive

Alkali Metals (Group 1)

• They lose the outer shell electron to form 1+ ions as you go down the group elements become more reactive
• This is because atoms get larger the outer shell electron is more easily lost.
• All alkali metals react with water -> alkali metal hydroxide + hydrogen gas.

Halogens (Group 7)

• They gain one electrons to form 1- ions.
• As you go down the group the elements get less reactive as the atoms get larger so the electron is more strongly attracted, therefore the incoming electron is less easily gained therefore the 1- ion is less readily. Vice Versa of Group 1
• In a displacement reaction, a more reactive element replaces a less reactive element in a compound.
• A more reactive halogen will displace a less reactive halogen from a compound. Halogen + metal -> metal Halide

Organic Chemsitry

• Organic chemistry definition: study of carbon compounds.
• Hydrocarbon definition: compound containing carbon and hydrogen only.
• Functional group: group of atoms giving characteristic chemical and physical properties
• Homologous series:
  • Same functional group, trend in physical properties, similar chemical properties and each one differing by CH2
• Isomers, Saturated, Unsaturated.

Alkenes

: Alkanes are saturated hydrocarbons: CnH2n+2

• Structure/Displayed Forumlas: Methane, Ethane, Propane, Butane, Pentane, Hexane

Alkane Reactions

• Only two, Combustions, and Substitution.
• Halogens, requires UV Light
• With lots of Oxygen, results in CO2 + H2O
• With a small oxygen supply, results in CO + C + H20

Alkenes:

Alkenes are unsaturated hydrocarbons with one C=C double bond

• General Formula CnH2n
• The chemical formula of But-2-ene and But-1-ene are the same, however, they have vastly different layouts.
  • The displayed formulas of two molecules can be isomers of each other.
  • Isomers are molecules that possess the same chemical formula.

Alkene Reactions

• Alkenes turn Bromine water from Orange to Colourless.
  • This is a test for C=C unsaturation
  • One product is created
  • This is also known as *Addition Reaction"

Monomer:

• A small molecule that can join with others to make a long chain polymer
• A polymer is a long chain of many monomers joined through Addition Polymerisation
• Requires one product.
• When a polymerisation occurs the double bond breaks and the chains lengthen
  • Addition polymers are usually saturated

1. Addition polymers are *inert"
2. Addition polymers do not biodegrade
3. When burned addition polymers create toxic gases

What is Crude Oil

1. Crude Oil: Mixture of Hydrocarbons with varying lengths, from small (CH4) to extremely large chains
   • Different fractions mean different intermolecular forces; this means, they all carry different boiling points

What is Fractional Distillation, and how does it work

• Method used to separate different boiling points in hydrocarbons
• Process
  • a Crude oil is heated to a high temperature and transferred into the fractionating column, and the tower has a temperature gradient (hot at the bottom, cool at the top
  • The different BPs of the hydrocarbons allow them to rise to their levels before either condensing or reaching the top
  • These Different levels are called “fractions"
• More molecules equal more viscous, less flammable, darker, and higher BPs.

Burning Fossil Fuel and its effect

1. Fossil Fuel: a fuel that derives from the fossilised remains of living organisms
   • Non-finite non-renewable resource.
2. Pollutants and formula
   • CO2 , NOx, SO2, C(soot)
3. What effects do these pollutants have on our environment
   • Global Warming + Ocean Acidification for SO2 and CO2
   • Respiratory issues

Cracking and when/why it happens

1. Cracking: the thermal decomposition of a long-chain alkane becomes shorter alkane molecules and an alkene
2. Cracking occurs when there are an excessive amount of larger non-flammable molecules this allows for smaller products for polymer to be made and burned.
3. Cracking requires high temperature (873k-973k) and a catalyst (Al2O3)"

Acids and Bases

• Aqueous solution of an acid: source of H+ ions

• If there isn't water, then there is no release of H+

• HCl(aq) => H2O + Cl-

• Base: substance that reacts with acids (OH-)

• Alkali: A base soluble in water with (OH- ions) in the aqueous solution

• NaOH + H2O => Na- + plus OH-

• Neutralisation: Acids and Alkalis cancel out to form water. H+ + OH- => H2O

• Acids a proton donors. A base is a proton acceptor. When an acid reacts, it transfers its H to the Base.

• Acid/base -> salt + water

• Acid + Metal -> Salt + H2

• Acid and Carbonate -> Salt + Carbon Dixiode + Water

• acid + amonia -> amonnia

• The pH Scale and Indicator and colour

• The lower the pH, the higher the concentration of H+

Solubility Rules

1. All sodium potassium,a nd AMMONIUM are SOLUABLE
2. All nitrates are soluble
3. Clorides r soluble excpet sliver and *lead"
4. sulfates are soluable except barium, calcium,*lead"
5. Carbonates are insoulable excepet potassium and ammonium
6. Hydroxies are *calcium * potassium and ammonium"

Preparation of Pure Salts with 3 methods

1. Titration Method- making a base from a soluble salt. (hcl and naoh). where you pipette acid.
2. Excess inoluable base- is making a base from a insoluble base -Measure out vol of acid

• Add salt. and warm the reaction but filter unreacted product.

3. Prepitation

• Insoluble salts from two soluable salts
• Mix soaliable and filtrat solid

Rates of Reactions

-Rate = Amount of product formed of reactant / Time

• Collision theory
• A reaction has to have enough energy to create a reaction due to the effects of the temperature, concentration and state. Factors affecting rates:
• Concentration - The higher the concentration, the more particles, and the frequency of higher speed collisions
• Catalyst- provides an alternate route to make a reaction faster and is not reused.

Experiments to measure rates

• Measuring the amount of gas produced
• Gas Syringe
• Measuring amount of mass using cotton wool
• Top hat balance
• See though rate
• Draw an X.