U2 AOS 1 Chapter 11

Question 1

half equation

Answer 1

reduction or oxidation equation of a complete redox reaction

Question 2

oxidation

Answer 2

species loses electrons

Question 3

oxidising agent

Answer 3

oxidant, chemical species that oxidises another substance by accepting electrons

Question 4

oxidation state

Answer 4

number assigned to an atom that can be used to determine the movement of electrons in a redox reaction, quantifies degree of oxidation based on counting electrons

Question 5

polyatomic ion

Answer 5

ions made of a group of atoms covalently bonded together

Question 6

redox reaction

Answer 6

A chemical reaction involving the transfer of one or more electrons between chemical species, oxidation and reduction occur simultaneously

Question 7

reduce

Answer 7

cause a chemical species to gain electrons

Question 8

reduction

Answer 8

gain of electrons

Question 8

reducing agent

Answer 8

reductant, chemical species that reduces another substance by donating electrons

Question 9

oxidation state rules

Answer 9

1. free elements (alone/bonded to itself) = 0 oxidation state eg. O2, C, Cl2
2. simple ion = oxidation state of the charge of the ion eg. Na+, O2-, Al3+
3. sum of oxidation states in a neutral compound = 0
4. sum of oxidation states in a polyatomic ion = charge of the ion eg. SO4^2-, NH4+
5. Hydrogen is +1 when bonded to nonmetals and -1 when bonded to metals.
6. Oxygen is -2
7. Transition elements’ oxidation numbers must be determined from the other elements in the compound.
8. The most electronegative atoms = negative oxidation state (gains electrons) eg. NO2, NH3

Question 10

identification of redox reactions

Answer 10

through the change in oxidation states (not necessarily change in state)

Question 11

oxidation states determine …

Answer 11

whether a redox reaction has taken place (shown by change), which species underwent reduction/oxidation

Question 12

metal displacement reaction

Answer 12

redox reaction in which a more reactive metal displaces a less reactive metal’s cation from solution eg. Fe(s) + Cu2+(aq) -> Fe2+(aq) + Cu(s)

Question 12

Balancing Redox Equations

Answer 12

KOHES,
1. key elements balanced

2. O balanced w/ water
3. H balanced w/ H+
4. electrons added to balance charge
5. states

Question 12

spontaneous redox reaction

Answer 12

readily occurring redox reaction that occurs w/out the addition of extra heat/energy

Question 12

reactivity series of metals

Answer 12

Organised scale of metals and their cations, ranked according to their strength as reducing and oxidising agents

Question 12

In a redox reaction producing H+, what happens to pH?

Answer 12

H+ on products not reactants, [H+ ions] increases, pH decreases

Question 13

why do OH groups oxidise easily?

Answer 13

Oxygen has 2 lone pairs, electrons can be easily lost (oxidise)

Question 14

reactivity series of metals trends

Answer 14

increasing reducing agent (tendency to oxidise), decreasing oxidising agent (tendency to reduce) from top to bottom; more reactive metals at the bottom

Question 15

higher reactivity (lower in the reactivity series)

Answer 15

stronger reducing agent, weaker oxidising agent

Question 16

most readily occuring spontaneous redox reactions

Answer 16

stronger oxidising and reducing agents

Question 17

lower reactivity (higher in the reactivity series)

Answer 17

weaker reducing agent, stronger oxidising agent

Question 17

galvanised

Answer 17

coated w/ zinc

Question 18

After a displacement reaction (Ag+Sn(s), the mass of the solid tin increased, why?

Answer 18

in displacement reaction Sn is lost from figurine forming Sn2+, however, Ag(s) is gained resulting in change in mass

Question 19

galvanic cell

Answer 19

electrochemical cell in which chemical energy from spontaneous redox reactions is converted into electrical energy

Question 20

anode

Answer 20

the negatively charged electrode at which oxidation occurs

Question 21

cathode

Answer 21

the positively charged electrode where reduction occurs

Question 22

electrode

Answer 22

chemically conductive medium

Question 23

sacrificial anode

Answer 23

A readily oxidised anode used to coat other metals for protection against corrosion

Question 24

sacrificial protection

Answer 24

The use of a protective layer that is readily oxidised

Question 25

dry corrosion

Answer 25

Oxidation of metals due to oxygen in the air

Question 26

wet corrosion

Answer 26

Oxidation of metals due to oxygen and water in the air

Question 27

rust

Answer 27

flaky reddish precipitate formed from the corrosion of iron w/ water and oxygen

Question 28

salt bridge

Answer 28

device used to seperate the solutions in each half cell while completing the circuit

Question 29

generating electricity in galvanic cell

Answer 29

electrons lost during oxidation flow through wire - provides electrical power

Question 30

direction of electrons in cell

Answer 30

towards positive cathode

Question 31

dry corrosion eg

Answer 31

aluminium - strong AlO3 layer forms on exposed metals, protects aluminium underneath from further corrosion

Question 32

wet corrosion eg.

Answer 32

1. Fe2+ + 2OH- -> Fe(OH)2 (iron hydroxide) 2. 4Fe(OH)2 + O2 + H2O -> 4Fe(OH)3 -> Fe2O3 x H2O (hydrated iron (III) oxide)

Question 33

preventing corrosion eg.

Answer 33

galvanised iron - iron coated w/ zinc used in steel for hull of ships to prevent corrosion, if Zn coating is damaged Zn still oxidises over iron

Question 34

chemical name and formula for rust

Answer 34

hydrated iron (III) oxide Fe2O3 x H2O

Question 35

freshwater vs seawater for iron corrosion

Answer 35

corrosion of iron requires oxygen and water

The reaction can be accelerated with a more electrically conductive solution

seawater is conducts electricity to a greater extent and contains more ions so corrosion is accelerated

Question 36

corrosion of aluminium vs iron

Answer 36

When oxidised, aluminium forms a protective layer of Al2O3, while iron forms rust which flakes easily

Al2O3 protects aluminium from further corrosion while rust does not

iron is more easily corroded than aluminium