Comprehensive Notes on Atomic Structure, Bonding, Fuels, and Metals

Atomic Structure & Bonding

• Nucleus: Contains protons and neutrons.

• Protons: Positive charge (p^+).

• Neutrons: No charge (neutral).

• Electrons: Orbit the nucleus with a negative charge (e^-).

• Neutral Atom: Number of protons equals the number of electrons.

• Atomic Number: Number of protons = number of electrons.

• Mass Number: Number of protons + number of neutrons.

• Isotope: Same atomic number (number of protons) but different mass number (number of neutrons).

  • Example: Sodium (Na) with atomic number 11 and mass number 23 or 25.

Covalent Bonding

• Occurs between non-metal atoms.

• Atoms share pairs of electrons to gain a full outer shell.

• Covalent Molecular: Small groups of atoms.

  • Low melting and boiling points.

  • Do not conduct electricity.

• Covalent Network: Large, continuous structure.

  • Very high melting points.

  • Usually do not conduct electricity (except graphite).

Ionic Bonding

• Occurs between metal and non-metal atoms.

• Ions are held together by strong electrostatic forces in a lattice structure.

• Ionic Compounds:

  • High melting and boiling points.

  • Conduct electricity when molten or in solution.

Catalyst

• Speeds up a reaction without being used up itself.

• Reduces the energy needed for the reaction.

Rates of Reaction

• As particle size decreases, the rate of reaction increases.

• As temperature increases, the rate of reaction increases.

• As concentration increases, the rate of reaction increases.

• Average Rate: Change in quantity / Change in time.

Chemical Formulae & Equations

• Chemical formulae show the types and numbers of atoms in a substance.

• Valency:

  • Group 1: Valency 1

  • Group 2: Valency 2

  • Group 3: Valency 3

  • Group 4: Valency 4

  • Group 5: Valency 3

  • Group 6: Valency 2

  • Group 7: Valency 1

  • Group 0: Valency 0

• Meaningful Names:

  • mon- 1

  • di- 2

  • tri- 3

  • tetra- 4

  • penta- 5

• Endings -ite and -ate contain oxygen.

  • Use charge, not valency.

• Diatomic Elements:

  • Hydrogen - H_2

  • Nitrogen - N_2

  • Oxygen - O_2

  • Chlorine - Cl_2

  • Fluorine - F_2

  • Bromine - Br_2

  • Iodine - I_2

• Spectator Ions: Do not take part in the reaction and occur in both reactants and products.

Titration

• Concentration (mol/L or M) = Number of moles / Volume (L)

Acids & Bases

• Acids: pH < 7 (litmus turns red).

  • Hydrochloric acid - HCl(aq)

  • Sulfuric acid - H2SO4(aq)

  • Nitric acid - HNO_3(aq)

  • Phosphoric acid - H3PO4(aq)

• Alkalis (Bases): pH > 7 (litmus turns blue).

• Soluble non-metal oxides dissolve in water forming acidic solutions.

• Dilute acids contain ions, therefore conduct electricity.

• Dilute acids contain hydrogen ions (H^+(aq)).

• During electrolysis, hydrogen ions are attracted to the negative electrode where they gain electrons to form hydrogen gas.

• Soluble metal oxides dissolve in water to form alkaline solutions.

• If a metal oxide is insoluble, it will not affect the pH of water.

• Neutralisation:

  • Acid + Base -> Salt + Water

  • Acid + Metal Hydroxide -> Salt + Water

  • Acid + Metal Oxide -> Salt + Water

  • Acid + Metal Carbonate -> Salt + Carbon Dioxide + Water

  • Acid + Metal -> Salt + Hydrogen Gas

Rates of Reaction (Revisited)

• High temperature: faster reaction.

• High concentration: faster reaction.

• Bigger particle size: slower reaction.

• Catalyst:

  • Speeds up the reaction without being used up itself.

  • Examples: Amylase (enzyme), Iron (in Haber process for ammonia production).

Atomic Structure to Bonding: Covalent Bonding

• Why do atoms form bonds? To achieve a stable outer energy level configuration.

• Forming Covalent Bonds:

  • Covalent bonds are between non-metal elements.

  • Covalent bonding involves sharing electrons.

• Formula for Covalent Compounds:

  • The valency of an atom is equal to the number of unpaired electrons.

• How do Covalent Bonds Hold Atoms Together?

  • When two atoms bond, the shared pair of electrons is attracted to the nuclei of both atoms.

• Covalent Molecules:

  • Only covalent compounds are made up of molecules.

  • Molecules are groups of atoms held together by covalent bonds.

• Covalent Networks:

  • Occur between a very large number of atoms.

Fuels

• Fossil fuels are formed from the remains of living organisms.

• Most common fossil fuels: Coal, oil, gas.

• Oil and gas are a group of compounds with similar boiling points, called a fraction.

  • Small molecules:

    • Evaporate easily

    • Very flammable

    • Flow easily

  • Large molecules:

    • Do not evaporate easily

    • Not as flammable

    • Do not flow easily

Structure of Hydrocarbons

• A hydrocarbon is a compound which contains carbon and hydrogen only.

• Methane: CH_4

• Ethane: C2H6

Homologous Series

• A homologous series is a family of compounds which have the same general formula and similar chemical properties.

• Alkanes General formula: CnH{2n+2}

• Alkenes General formula: CnH{2n}

• Cycloalkanes General formula: CnH{2n}

• The complete combustion of a hydrocarbon produces carbon dioxide and water.

• Combustion gives out energy.

• Any reaction or process which gives out energy is known as exothermic.

• Any reaction or process which takes in energy is known as endothermic.

Energy from Fuels

• Units for energy are joules (J) or kilojoules (kJ).

• Specific heat capacity of water: 4.18 kJ/kg/°C

• E = mc\Delta T

  • E: Energy gained by the water (J)

  • m: Mass of water heated (kg) (1 ml = 1 cm³ = 1 g; 100 ml = 100 cm³ = 100 g = 0.1 kg)

  • c: Specific heat capacity of water (4.18 kJ/kg/°C)

  • \Delta T: Change in temperature of water (°C)

• Problems with Energy Release Experiment:

  • Lots of heat energy is lost to the surrounding air and to the apparatus.

  • Alcohol is burning with a yellow sooty flame, which means it is not burning efficiently (incomplete combustion).

Alkanes, Alkenes and Cycloalkanes

• Alkanes: End in -ane. Saturated hydrocarbons with only single bonds between the carbon atoms. General formula: CnH{2n+2}. Example: Methane (CH4), Ethane (C2H_6).

• Alkenes: End in -ene. Unsaturated hydrocarbons with double bonds between the carbon atoms. General formula: CnH{2n}.

• Cycloalkanes: Saturated hydrocarbons with carbon atoms arranged in a ring. General formula: CnH{2n}.

• Naming Convention:

  • meth- 1 carbon

  • eth- 2 carbons

  • prop- 3 carbons

  • but- 4 carbons

  • pent- 5 carbons

  • hex- 6 carbons

  • hept- 7 carbons

  • oct- 8 carbons

  • non- 9 carbons

  • dec- 10 carbons

• Alkanes do not decolourise bromine water.

• Alkenes do decolourise bromine water (addition reaction).

• Hydrogenation: Adding hydrogen to an alkene to make an alkane.

• Hydration: Adding water to an alkene to make an alcohol.

• Isomers: Molecules with the same molecular formula but different structural formulas.

• Alcohols: Contain a hydroxyl (-OH) functional group and end in -ol. Example: Ethanol (C2H5OH).

• Carboxylic Acids: Contain a carboxyl group (-COOH). Example: Ethanoic acid (CH_3COOH).

• Esters: Formed when an alcohol and a carboxylic acid react.

Metals & Alloys

• Metal + Oxygen -> Metal Oxide

• Metal + Water -> Metal Hydroxide + Hydrogen Gas

• Acid + Metal -> Salt + Hydrogen Gas

• The name of the salt depends on the acid and metal.

  • Hydrochloric acid -> Metal Chlorides

  • Sulfuric acid -> Metal Sulfates

  • Nitric acid -> Metal Nitrates

  • Phosphoric acid -> Metal Phosphates

Ion-Electron Half Equations

• OILRIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).

• Oxidation example: Na -> Na^+ + e^-

• Reduction example: Cl_2 + 2e^- -> 2Cl^-

Reactivity Series of Metals

• Potassium (K)

• Sodium (Na)

• Lithium (Li)

• Calcium (Ca)

• Magnesium (Mg)

• Aluminum (Al)

• Zinc (Zn)

• Iron (Fe)

• Lead (Pb)

• Copper (Cu)

• Mercury (Hg)

• Silver (Ag)

• Gold (Au)

• Metals higher in the series react with cold water.

• Metals further down react with dilute acids.

• The most reactive metals react with oxygen easily.

Metal Ores

• Ores are naturally occurring rocks that contain metals or metal compounds in sufficient amounts to be worth extracting.

Reactivity & Reduction

• Removing a metal from its oxide.

• If a metal is very reactive, it forms strong bonds, requiring significant energy for extraction (so the metal oxide seems unreactive).

• If a metal is very unreactive, it forms weak bonds, requiring little energy for extraction (so the metal oxide seems reactive).

• A reducing agent is a substance which causes another substance to be reduced (supplies electrons).

Smelting

• A process for extracting a metal from its ore by mixing it with carbon and heating it (industrial scale).

Percentage Composition

• \% \text{ composition} = \frac{\text{mass of element}}{\text{formula mass}} \times 100\%\newline Formula mass:

• Given Water (H_2O)

  • H atomic mass = 1

  • O atomic mass = 16

  • (2 \times 1) + (1 x 16) = 18

• You can then calculate the mass of each element:

  • Total mass of compound. (Examples: C6H12O6)

    • (6 x 12) + (12 x 1) + (6 x 16) = 180

  • mass oxygen\frac{(6 \times 16)}{180} \times 100 = \frac{96}{180} \times 100

Making Electricity

• A cell is an apparatus which generates an electrical current from a chemical reaction.

• A battery is two or more cells connected together.

• Electrochemical Series:

  • The higher a metal is in the electrochemical series, the easier it is to lose electrons (form ions).

  • When metals are arranged in decreasing voltage, they form an electrochemical series.

• In a cell:

  • Electrons flow from the metal higher in electrochemical series through the wire to the metal lower in the series.

  • The further apart the metals are in the series, the higher the voltage produced.

• An electrolyte is a substance, usually an ionic solution, which conducts electricity.

• Ion Bridge: Allows only ions to flow across the bridge, not electrons.

• Electrons flow in wires only.

Redox Reaction

• Involves a transfer of electrons.

• Look for:

  • Ion changing to atom.

  • Atom changing to ion.

  • Size of charge changes.