Chemistry Unit 13

Self Ionisation of Water

-If the concentration of H3O+ ions increases due to the presence of a dissolved acid, the concentration of OH− ions decreases such that the product is still 10−14. The reverse applies if a source of OH− ions (a dissolved base) is present. Therefore, for any acidic, basic or neutral solution, we may write: Kw = [H3O+][OH−] = 10−14 -Kw is termed the self-ionisation constant for water, and [H3O+][OH−] is called the ionic product. -If [H3O+] > [OH−], the solution is acidic. -If [OH−] > [H3O+], the solution is basic. -If [H3O+] = [OH−], the solution is neutral.

The pH meter

-Most laboratories in industry have a pH meter that is used to make rapid, accurate measurements of pH. It can be connected to a computer to monitor pH changes continuously. -The voltage of the electrode changes with the [H3O+] in the solution into which it is dipped. Values of pH obtained by a pH meter are accurate to within 0.01 units of the true pH and are not affected by the colour and cloudiness of the unknown solution.

Acid rain

-Rain containing acids that form in the atmosphere when industrial gas emissions (especially sulfur dioxide, nitrogen oxides and also carbon dioxide) combine with water. -Rain that has a pH lower than 5.

Strengths of Acids and Bases

-Strong acids or bases donate or accept protons readily. -Weak acids or bases do not donate or accept protons readily. -A strong acid donates protons readily and forms a large number of ions in water. Strong acids form solutions that are good conductors of electricity. Weak acids do not donate protons readily. As a result, only a few ions are formed by reaction with water and the resulting solutions are poor electrical conductors. -The strength of a base also affects its conductivity. A base is said to be strong if it produces many hydroxide ions in solution because the hydroxide ion readily accepts protons. However, a weak base such as ammonia does not readily ionise and accept protons. It forms very few ions in solution. Ammonia is therefore a poor conductor of electricity. -Completely ionise=→ -Partially ionise=⇌

Strength versus concentration

-The strength of an acid is different from the concentration of an acid. The strength of a solution is determined by the number of ions present. -Concentration refers to the amount of an acid or base that is dissolved in a given volume of water. A large amount always produces a concentrated solution whereas a small amount in the same volume of water produces a dilute solution. It is possible to have a weak, concentrated acid or a dilute solution of a strong acid.

Neutralisation reaction

A chemical reaction in which an acid and a base react together to produce a salt and water. -Since Na+ and Cl− are spectator ions and do not take part in the reaction, neutralisation may be shown simply as: H+(aq) + OH−(aq) → H2O(l), this is known as the ionic equation of neutralisation. -The hydrogen ion, H+, does not exist by itself in water. Neutralisation reactions are one way of producing pure samples of salts. The water can be removed by evaporation.

Steps for balancing half-equations

1. The half-equations for oxidation and reduction showing conjugate pairs. 2. Balance all elements except hydrogen and oxygen. 3. Balance oxygen atoms, where needed, by adding water. 4. Balance hydrogen atoms, where needed, by adding H+. 5. Balance the charge by adding electrons. 6. Multiply each half-equation by factors that will lead to the same number of electrons in each half-equation. 7. Add the half-equations and omit the electrons.

Reactions of Acids

1. acid + metal (not Cu, Hg or Ag) → salt + hydrogen gas H2SO4(aq) + Zn(s) → ZnSO4(aq) + H2(g) 2H+(aq) + Zn(s) → Zn2+(aq) + H2(g) 2. acid + metal carbonate → salt + carbon dioxide gas + water H2SO4(aq) + CaCO3(s) → CaSO4(aq) + CO2(g) + H2O(l) 2H+(aq) + CaCO3(s) → Ca2+(aq) + CO2(g) + H2O(l) 3. acid + metal hydrogen carbonate → salt + carbon dioxide gas + water HCl(aq) + NaHCO3(s) → NaCl(aq) + CO2(g) + H2O(l) H+(aq) + NaHCO3(s) → Na+(aq) + CO2(g) + H2O(l) 4. acid + metal sulfite → salt + sulfur dioxide gas + water 2HCl(aq) + FeSO3(aq) → FeCl2(aq) + SO2(g) + H2O(l) 2H+(aq) + SO32−(aq) → SO2(g) + H2O(l) 5. acid + metal sulfide → salt + hydrogen sulfide gas 2HCl(aq) + FeS(s) → FeCl2(aq) + H2S(g) 2H+(aq) + FeS(s) → Fe2+(aq) + H2S(g) 6. acid + metal oxide (basic oxide) → salt + water 2HNO3(aq) + MgO(s) → Mg(NO3)2(aq) + H2O(l) 2H+(aq) + MgO(s) → Mg2+(aq) + H2O(l) 7. acid + base (metal hydroxide) → salt + water H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l) H+(aq) + OH−(aq) → H2O(l)

Reduction

A decrease in the oxidation number; a gain of electrons. -Oxidant/ Oxidising agent is the species that is reduced.

Triprotic acid

An acid able to donate three protons per molecule.

Monoprotic acid

An acid that can donate only one proton (hydrogen ion) per molecule.

Diprotic acid

An acid that can donate two protons per molecule

Oxidation

An increase in the oxidation number; a loss of one or more electrons. -Reductant/Reducing agent is the species that is oxidised.

Half-equations

Chemical equation showing either oxidation or reduction in a redox reaction

Displacement reactions

Chemical reaction in which a more reactive element displaces a less reactive element from its compound. -Aluminium metal often appears to be less reactive than the reactivity series indicates because it has a coating of aluminium oxide and this protects it from reacting further with oxygen. This coating can be thickened by a process called anodising.

Acids

Compounds that form hydrogen ions when dissolved in water. -A substance that can donate a proton to a base. Acids have many common properties. They: -Usually taste sour. -Are corrosive. -Are molecular in structure and dissolve in water to produce an electrolyte (substance that conducts electricity). -Affect the color of certain natural and synthetic dyes (they turn litmus, a plant dye, from blue to red). -Are neutralized by bases.

Bases

Compounds that reduce the concentration of hydrogen ions in a solution. -A substance with a pH greater than 7. -Compounds with properties that in some ways complement those of acids. Bases have many common properties. They: -Usually taste bitter. -Feel slippery (they react with the natural oils in the skin and produce soap, which gives the characteristic 'slippery feel'). -Turn litmus from red to blue are electrolytes (substances that conduct electricity). -May be corrosive (for example, NaOH, which breaks down organic substances such as fat, hair and vegetable matter) and are good drain cleaners. -Are generally ionic substances. -Are oxides or hydroxides of metals (ammonium hydroxide is a base). -Are usually insoluble in water. A base that is soluble in water is called an alkali.

Oxidation-reduction/Redox reaction

Describes a reaction that involves the transfer of one or more electrons between chemical species. -Oxidation cannot occur unless reduction occurs simultaneously. -When writing equations for redox reactions, we do not show electrons as all electrons given off during oxidation are taken in during reduction. However, if we are considering oxidation or reduction reactions separately, it is appropriate (and necessary) to write reactions that do show the electrons. Such equations are called partial equations or half-equations.

Dilution

Dilution is the process of adding water to a solution. For solutions that are acidic or basic, this affects the concentration of H3O+ ions that are present and, hence, the pH. Because pH is on a log scale, the factor by which H3O+ changes is not the same as that by which pH changes.

The production of the Hydronium Ion

HCl(g) + H2O(l) → H3O+(aq) + Cl−(aq) -Along with the chloride ion, a hydronium ion is produced. A proton, H+, cannot exist by itself as it is attracted to the negative end of the polar water molecule to form a hydronium ion, H3O+.

Ionisation reactions

Ionisation is a reaction in which a substance reacts with water to produce ions, OH- or H30+

spectator ions (net ionic equations)

Just hanging, not really doing anything, remains in the solution unchanged, can ignore this. -Not in Ionic equation of neutralisation.

Polyprotic acids

Polyprotic acid can donate more than one proton per molecule of acid.

Amphiprotic substances

Some substances can act as acids or bases, according to their chemical environment. This means that they can donate or accept protons. Such substances are described as amphiprotic. -Ampholytes are ionic amphiprotic substances.

Indicators

Substances that can indicate an acid or basic state and a range of pH values

The Brønsted-Lowry theory of acids and bases

The Brønsted-Lowry theory defines acids and bases as follows: -An acid is a substance that donates a proton (H+ ion) to another substance. -A base is a substance that accepts a proton (H+ ion) from another substance. Since a hydrogen atom is simply a proton and an electron, removing the electron leaves a proton, H+. -This means that a reaction between an acid and a base involves a proton transfer from the acid to the base. This occurs only when both an acid and a base are present. -The proton described by Brønsted and Lowry is simply a hydrogen ion, H+. The hydrogen ion is transferred from one substance to another. The substance that loses an H+ ion is the acid and the one that accepts it is the base.

Acids and Bases

The characteristic properties of an acid can be reduced by adding a base. If enough base is added, the acidic properties completely disappear. When this occurs, we say that the acid has been neutralized by the base.

The pH scale

The pH scale is usually applied over a range from 1 to 14. -Using this scale, a neutral solution has a pH of 7. -Values lower than 7 indicate an acidic solution. -Values higher than 7 indicate a basic/alkalinic solution.

Oxidant/Oxidising agent

The species that causes oxidation of another substance and is itself reduced. -Electron acceptor.

Reductant/Reduing agent

The species that causes reduction of another substance and is itself oxidised. -Electron donor.

Conjugate redox pair

Two species that differ only by a certain number of electrons

Conjugate acid-base pairs

When an acid and a base react, a conjugate acid and base are formed. -In an acid-base reaction, the substance acting as the acid gives away a proton and forms a conjugate base. The substance acting as a base, after accepting a proton, forms its conjugate acid. An acid-base reaction therefore forms two conjugate pairs, with the formulas of each pair member differing by the H+ ion that was transferred. -Conjugate base = acid − H+. -Conjugate acid = base + H+. -Conjugate acid-base pairs differ by a proton, H+. To find the conjugate base of an acid, we subtract one H+; to find the conjugate acid of a base we add one H+.

Hydrolysis(a type of ionisation reaction)

When an ionic substance dissolves in water, the resulting solution is often not neutral. This is because the ions can act as acids or bases when reacting with water and produce solutions that are either acidic or basic; this is known as hydrolysis. A hydrolysis reaction is one in which a substance reacts with water to form OH− or H3O+ ions. -Solution acidic if H3O+ is formed. -Solution alkline(basic) if OH- is formed.

Dissociation of Bases

When ionic bases dissolve in water, they dissociate or separate into their constituent ions. They do not ionise since they do not actually react with the water to produce ions as acids do.

Universal indicators

a mixture of separate indicators that show colour changes in solutions of a large range of pH values

pH formula

pH=−log10[H3O+] -Square brackets are used to denote concentrations measured in M. [H3O+] = 10−pH