Hein's Chem Flashcards

Electrochemistry SL

Main reactions in chemistry

Main reactions in chemistry

• acid-base reactions

• precipitation reactions

• redox reactions

redox reaction

involves two processes, reduction and oxidation

Ways of considering reduction and oxidation

• in terms of specific elements - oxygen and hydrogen

• in terms of electron transfer

• in terms of oxidation state

Oxidation (combining with oxygen)

at the simplest level oxidation can be considered as a reaction in which a substance combines with oxygen

Oxidation (combining with oxygen) examples

Reduction (Removal of oxygen)

a reaction in which a substance removes/losses its oxygen

Reduction (Removal of oxygen) example

Nickel (II) Oxide is reduced (loses O)

Reduction (addition of hydrogen)

a reaction in which a substance adds/gains hydrogen

Reduction (addition of hydrogen) example

Hydrogen is added, thus oxygen is removed from the tungsten (VI) oxide

Oxidation (electron transfer)

losses electrons

Reduction (electron transfer)

gains electrons

OILRIG

Oxidation (electron transfer) example

Reduction (electron transfer) example

Why can't all redox processes be considered in term of redox processes?

Carbon dioxide is molecular, with covalent bonds, so no ionic bonds form and therefore can't be explained in terms of electron transfer as in theory no electrons are lost or gained and the carbon dioxide is a neutral species. Making the original definition of the addition of oxygen more appropriate here

Oxidation (Oxidation state)

a process in which the oxidation state increases

Reduction (Oxidation state)

a process in which the oxidation state decreases

Oxidizing agents (Oxidants)

substances that cause another species to be oxidized by itself reducing by gaining an electron and decreasing in oxidation state

Reducing Agents (Reductants

substances that cause another species to be reduced by itself oxidizing by donating/losing an electron and increasing in oxidation state

Rules of assigning oxidation state

1. Elements only combined with them self = 0 (O2, Na, C60)

2. Alkali metals (group 1) are always +1

3. Alkali earth metals (group 2) are always 2

4. Aluminium is usually +3

5. Hydrogen is +1 (bonded to non-metals), except for metal hydrides (-1), like NaH

6. Oxygen is -2 except for F2O (+2) and peroxides (-1)

7. Fluorine is always -1

8. Group 17 is usually -1, except when with F and O (in HClO4, Cl = +7

9. Sum of oxidation numbers = 0 for neutral molecules, or will match the overall change for ions

Variable Oxidation States

The occurrence of variable oxidation states and, often, the interconversion between them is a characteristic of most d- block metals (especially transition metals)

Type A transition metals

• Sc, Ti, and V

• FIND

Type B transition metals

Cr and Mn

Type C transition metals

Fe, Co, Ni, Cu, and Zn

Stock nomenclature System

Roman numerals (I, II, III etc) are used to indicate the oxidation number

Naming oxoanions

• If there is only one oxoanion (-ate)

• If there are two oxoanions the smaller number of oxygen is (-ite) and larger number of oxygen is (-ate)

• If there are four oxanions the smaller number of oxygens will end (-ite) and be prefixed by (hypo), next will end in (-ite), the thid will end in (-ate) and the largest number of oxygens will end in (-ate) and be prefixed by (per)

Half Equations

Equations describing one of either the oxidation or reduction processes of a redox reaction.

Working Method (Half Equations)

Step 1: Assign oxidation states for each atom in the reactant and product species. Step 2: Deduce which species is oxidized and which species is reduced Step 3: State the half-equations for the oxidation process and the corresponding half-equation for the reduction process Step 4: Balance these half-equations so that the number of electrons lost equals the number of electrons gained Step 5: Add the two half-equations together to write the overall redox reaction Step 6: Check the total charge on the reactant and product sides Step 7: Balance the charge by adding H+ and H20 to the appropriate sides

Activity Series

Ranks metals according to the ease with which they undergo oxidation.

• More reactive and greater ease of oxidation are at the top

• less reactive and less ease of oxidation are at the bottom

Metals higher in the reactivity series can...

displace those lower down from solutions of their respective salts

Activity Series Data booklet

section 25

Group 17 reactivity series

Activity Series Example

• Cl has smaller atomic radii so is more electronegative and therefore has a greater attraction for an electron than Br so Cl takes its electron/place

• Reaction seen as a color change from clear to Yellow/Orange due to the formation of aqueous bromine

Chlorine in Water disinfection

• cannot be used to treat viruses

• leaves a residual taste and unpleasant odor

• can form toxic by-products, often carcinogenic

• Cheaper

How is Chlorine used to disinfect water

added chlorine gas (Cl2), sodium hypochlorite (NaOCl), or calcium hypochlorite, Ca(OCl)2 which all yeild hypochlorous acid (HOCl) which is antibacterial

Ozone in Water Disinfection

• can be used to treat viruses

• leaves no unpleasant residual taste or odor

• fewer toxic by-products

• more expensive

Redox Titration Formulae used in volumetric analysis

Redox Titration Working Method (Textbook)

Step 1: Deduce the balanced redox equation (using oxidation states) Step 2: Identify values for 3 of the 4 possible given data (Va, ca, Vb, and cb) and stoichiometry coefficients (va and vb) Step 3: Set up expression and fill in the known data (see image) Step 4: Solve for the missing data (Va, ca, Vb, or cb) Step 5: Check that value found is the same as the one the equation asks for

Redox Titration Working Method (website)

Regular titration calculations Step 1: Determine the overall balanced redox equation Step 2: Determine the amount of titration solution needed in the volume of the original Step 3: Use the balanced equation to determine the amount of Fe2+ present initially Step 4: Determine the concentration in mol dm^-3

Disproportionation

when an element is both oxidized and reduced in the same reaction

Concentration in parts per million

Units of ppm

mg/L or mg dm^-3

Solubility of Oxygen in water

low

Temperature affect on the solubility of oxygen in water

as temperature increases, the solubility of water decreases

Biochemical Oxygen Demand (BOD)

the amount of oxygen required to oxidize organic matter in a sample of water at a definite temperature over a period of 5 days (Measured in units of ppm)

Maximum amount of Lead(II) cations allowed in drinking water according to the WHO

0.001 mg dm^-3 or 0.001 ppm

BOD of less than 1 ppm

Pure water

BOD of 20 or more

poor quality

BOD examples

Winkler Method

Step 1: Add an excess amount of manganese(II) so oxygen can oxidize the Mn(II) to make Mn(VI) ions Step 2: Add Potassium Iodine which will be oxidized by the manganese(VI) salt to form iodine (I2) Step 3: Iodine is titrated with standard sodium thiosulfate solution

Molar Ratio in the Winkler Method

For every one mole of dissolved oxygen, four moles of thiosulfate ions are needed

Water Quality Chart

Energy

the capacity to do work measured in Joules (J)

Law of conservation of energy

Energy cannot be created or destroyed but is converted from one form to another

Voltaic (or galvanic) cells

convert chemical energy (from spontaneous exothermic chemical processes) to electrical energy

Conversion in the Voltaic Cell half equation

Electrolytic Cells

convert electrical energy to chemical energy, by bringing about a non-spontaneous process

Electrodes

a conductor of electricity used to make contact with a non-metallic part of a circuit

Electrolyte

solution in a cell

Types of Electrodes

• anode is positive

• cathode is negative

Polarity in a voltaic cell

• cathode is the positive electrode

• anode is the negative electrode

Polarity in an electrolytic cell

• cathode is the negative electrode

• anode is the positive electrode

PANIC

CROA

Oxidation (voltaic and electrolytic cells)

oxidation always takes place at the anode

Reduction (voltaic and electrolytic cells)

reduction always takes place at the cathode

Types of Electrodes In a voltaic cell

• Metal/metal-ion

• Metal ions in two different oxidation states

• gas-ion

voltaic cell (Metal/metal-ion )

consists of a bar of metal dipped into a solution containing cations of the same metal which is separate from another half-cell

Examples of voltaic cell (Metal/metal-ion )

• Fe(s)|Fe^2+(aq)

• Zn(s)|Zn^2+(aq)

• Cu(s)|Cu^2+(aq)

How are two voltaic cells (Metal/metal-ion ) connected

Liquid junction called a salt bridge

What is a salt bridge?

a concentrated solution of a strong electrolyte which allows ions to diffuse out of it

Salt Bridge Function

• allows for the physical separation of the cathode (reduction process) and anode (oxidation process), preventing mixing of the two solutions

• provides electrical continuity (path for the migration of the positive ions (cations) and the negative ions (anions)

• reduces liquid-junction potential (voltage created by the two solutions)

liquid junction potential

the voltage generated when two different solutions come into contact with each other, which occurs due to an unequal cation and anion migration across the junction

Daniell Voltaic Cell

• Zn(s)|Zn^2+(aq) (Oxidation at anode)

• Cu(s)|Cu^2+(aq) (Reduction at cathode)

What happens in the Daniell Voltaic Cell

• Blue color of the copper (II) fades

• copper bar increases in size as it is coated in more copper

• zinc bar gets thinner as zinc ions are lost

Regular Salt bridges (Especially Daniell Voltaic Cell)

• Sodium sulfate (Na2SO4)

• Potassium Chloride (KCl)

How to draw a voltaic cell

Cathode is drawn on the right-hand side

How to determine which metal will be oxidized and reduced in a voltaic cell?

the metal which is the highest in the reactivity series is oxidized

Cell Diagrams

a convenient shorthand to represent a voltaic cell

• the anode written on the left and the cathode on the right

• Salt bridge represented by two parallel lines

• follow RO||OR

Electrolysis Definition

the process by which electrical energy is used to drive a non-spontaneous chemical reaction

Electrolytic Cell Contains

• single container

• two electrodes (cathode and anode)

• electrolyte (a solution)

• battery

Electrolysis Example (Molten Salt of lead(II) bromide)

• anode (positive electrode): oxidation 2Br → Br2(g) + 2e

• cathode (negative electrode): reduction Pb2(l) + 2e → Pb(l)

• overall cell reaction: PbBr2(l) → Pb(l) + Br2(g)

Electrolysis of molten salt (working method)

Step 1: identify all species present Step 2: identify which species are attracted to the cathode and which are attracted to the anode Step 3: Deduce the two half-equations taking place at the cathode and anode and thee overall cell reaction Step 4: Draw and annotate the electrolytic cell and show the direction of movement of electrons and the direction of ion flow Step 5: State what would be observed at each electrode