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U1&2 Chemistry > U2 AOS 1 Chapter 11 > Flashcards

U2 AOS 1 Chapter 11 Flashcards

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1
Q

half equation

A

reduction or oxidation equation of a complete redox reaction

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2
Q

oxidation

A

species loses electrons

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3
Q

oxidising agent

A

oxidant, chemical species that oxidises another substance by accepting electrons

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4
Q

oxidation state

A

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

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5
Q

polyatomic ion

A

ions made of a group of atoms covalently bonded together

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6
Q

redox reaction

A

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

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7
Q

reduce

A

cause a chemical species to gain electrons

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8
Q

reduction

A

gain of electrons

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8
Q

reducing agent

A

reductant, chemical species that reduces another substance by donating electrons

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9
Q

oxidation state rules

A
  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
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10
Q

identification of redox reactions

A

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

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11
Q

oxidation states determine …

A

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

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12
Q

metal displacement reaction

A

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)

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12
Q

Balancing Redox Equations

A

KOHES,
1. key elements balanced

  1. O balanced w/ water
  2. H balanced w/ H+
  3. electrons added to balance charge
  4. states
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12
Q

spontaneous redox reaction

A

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

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12
Q

reactivity series of metals

A

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

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12
Q

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

A

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

13
Q

why do OH groups oxidise easily?

A

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

14
Q

reactivity series of metals trends

A

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

15
Q

higher reactivity (lower in the reactivity series)

A

stronger reducing agent, weaker oxidising agent

16
Q

most readily occuring spontaneous redox reactions

A

stronger oxidising and reducing agents

17
Q

lower reactivity (higher in the reactivity series)

A

weaker reducing agent, stronger oxidising agent

17
Q

galvanised

A

coated w/ zinc

18
Q

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

A

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

19
Q

galvanic cell

A

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

20
Q

anode

A

the negatively charged electrode at which oxidation occurs

21
Q

cathode

A

the positively charged electrode where reduction occurs

22
Q

electrode

A

chemically conductive medium

23
Q

sacrificial anode

A

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

24
Q

sacrificial protection

A

The use of a protective layer that is readily oxidised

25
Q

dry corrosion

A

Oxidation of metals due to oxygen in the air

26
Q

wet corrosion

A

Oxidation of metals due to oxygen and water in the air

27
Q

rust

A

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

28
Q

salt bridge

A

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

29
Q

generating electricity in galvanic cell

A

electrons lost during oxidation flow through wire - provides electrical power

30
Q

direction of electrons in cell

A

towards positive cathode

31
Q

dry corrosion eg

A

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

32
Q

wet corrosion eg.

A
  1. Fe2+ + 2OH- -> Fe(OH)2 (iron hydroxide) 2. 4Fe(OH)2 + O2 + H2O -> 4Fe(OH)3 -> Fe2O3 x H2O (hydrated iron (III) oxide)
33
Q

preventing corrosion eg.

A

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

34
Q

chemical name and formula for rust

A

hydrated iron (III) oxide Fe2O3 x H2O

35
Q

freshwater vs seawater for iron corrosion

A

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

36
Q

corrosion of aluminium vs iron

A

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

U1&2 Chemistry (10 decks)

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