Year 10 Chemistry: Chemical Reactions

• Chemical reactions explained for Year 10 Chemistry students

Nanoparticles

• Nanoparticles have a diameter between 1 and 100nm.
• Nanoparticles are bigger than a molecule (around 0.5nm) or an atom (around 0.1nm).
• They are smaller than a cell, which is around 7000nm.
• Nanometre (nm) equals 1x10⁻⁹.
• An atom is approximately 1x10⁻¹⁰.
• Nanoparticles have a high surface area to volume ratio.
• Nanoparticles may have unknown long-term effects and can pass through bodies.
• Particles of this size behave differently compared to the same chemical in bulk because of their very high surface area to volume ratio.

Conservation of Mass

• Products and reactants of an equation must have the same total number of each atom.
• Atoms cannot be created nor destroyed.
• Mass is conserved in any reaction.

Ionic Equations

• Many reactions in solution only make sense when considering the ions.
• Ionic equations focus on the substances that undergo change.
• Ionic equations will always show ions and contain charges.
• Spectator ions, which do not change state or charge, are not written.
• Soluble ionic compounds are written as their ions.
• Insoluble ionic substances are written as solids.
• Metals and covalent compounds are written in full because they are not made of ions.
• Precipitation reactions occur when a solid is formed by two soluble ions (aq) reacting together.

Ionic Half Equations

• Ionic half equations show where electrons are added or taken away in a redox reaction.
• Balancing half equations involves balancing the number of atoms, working out the overall charge on each side, identifying the more positive side, and adding electrons to the more positive side to balance the charge.

pH Scale

• pH measures how acidic or alkaline a solution is.
• Water is neutral and has a pH of 7.
• Lower pH values indicate more acidic solutions, while higher values indicate more alkaline solutions.
• pH can be tested using a universal indicator paper, universal indicator solution (UI), or a pH meter.
• Universal Indicator results include, Vimto pH of 1.5, Tap Water with pH of 7.5, Sodium Hydrogen Carbonate with pH of 11, Dishwater Tablet with pH of 12, Blueberry Lemonade with pH of 2, Fizzy Drink with pH of 2.5

PH Meter

• To accurately measure pH with a pH meter, wash the probe with water and then calibrate it using a calibration buffer, adjusting to match the pH of the buffer solution.
• Some meters may need calibration against two or three buffers.

pH Calculations

• Type 1: Determine a pH from concentration of H+ ions
• Count the number of zeroes in the concentration to find the pH.
• For a solution containing 0.0001 mol/dm³ H+, the pH is 4.
• For a solution containing 1 x 10⁻⁹ mol/dm³ H+, the pH is 10.
• Type 2: Changing the concentration by a factor of 10 will change the pH
• For each factor of 10 that a solution gets more dilute, the pH increases by 1.
• If a solution with a pH of 3.1 is diluted by a factor of 1000, its pH becomes 6.1 (3.1 + 3 = 6.1).
• If a solution of acid has a pH of 2.83 and is concentrated 100 times, its pH becomes 0.83 (2.83 - 2 = 0.83).

Moles

• It is not practical in chemistry to work with one atom, so moles are therefore used.

Using Formulas and Equations

• A formula gives the ratio of chemicals in a compound.
• A balanced equation gives the ratio of moles of chemicals in a reaction.
• Process involves writing a balanced equation, calculating all molar masses needed, calculating moles, adjusting the chemical scale as needed and solving.

Energy Profile Diagrams and Reactions

• Reaction energy diagrams show what happens to energy when bonds are broken and remade during a reaction.
• Breaking bonds is endothermic and requires energy.
• Making bonds is also endothermic and releases energy.
• The overall reaction is exothermic/endothermic where Overall, energy is determined by which is dominant.

Catalyst

• A catalyst provides an alternative pathway for the same reaction with less energy.

Redox Reactions

• Oxidation involves adding oxygen or removing hydrogen.
• Reduction involves removing oxygen or adding hydrogen.
• A redox reaction involves both reduction and oxidation.
• Electrons are transferred from one atom to another atom.
• Oxidation is the loss of electrons.
• Reduction is the gain of electrons.
• If one atom is losing electrons, another must gain electrons.

Displacement Reactions

• A more reactive metal can replace/displace a less reactive metal.
• Magnesium, Zinc, Iron and Copper show the reactivity order with Magnesium being most reactive.

Acids

• Acids all have a pH less than 7 and release hydrogen ions (H⁺) when reacting or can all react the same way.
• Strong acids completely ionize in water.
• Weak acids do not fully ionize in water.
• Examples of the three strong acids: Hydrochloric acid (HCl) makes Chlorides, Nitric acid (HNO₃) makes Nitrates, and Sulphuric acid (H₂SO₄) makes Sulphates.
• Neutralization is possible for all acids
• Ethanoic acid (vinegar) is an example of a weak acid..
• Strong acids have a lower pH (1-3) compared to weak acids (4-6) at that same concentration.
• Strong acids react faster than weak acids and the same amount of product is eventually made.
• Concentration indicates the number of moles of acid in a dm³ of solution, and strength indicates how fully it ionizes.