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Chemistry > 9. Metals > Flashcards

9. Metals Flashcards

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1
Q

Physical properties of metals

A

Electrical conductivity - Metals are good conductors bc of delocalised electrons formed during metallic bonding. (non-metals are good insulators). (Wires are made of conductive metal and coated with non-metal insulators)

Thermal conductivity - Metals allow heat to easily pass through them

Melting and Boiling points - Most metals have high melting points and extrememly high boiling points.

Malleability - Metals are easily bent + shaped. When heated they become even more malleable as the atoms have mroe energy and can move more easily.

Ductility - They can be stretched into wires. This is because the atoms easily reorganise themselves.

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2
Q

Chemical properties of metals (list)

A

Reactions with dilute acids…

Reactions with water…

Reactions with oxygen…

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3
Q

Metal reacting with dilute acid (eg. zinc + hydrochloric acid)

(NEED TO KNOW REACTION OF ZINC WITH ACID)

A

metal + acid → salt + hydrogen

(Salt = compound, not actual salt)

Observation - Metal slowly disappears (turning into a soluble salt that dissolves in the water), Hydrogen gas bubbles up (Bubble slowly form).

zinc + hydrochloric acid –> zinc chloride + hydrogen

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4
Q

Metal reacting with water (eg. Calcium + water)

(NEED TO KNOW REACTION OF CALCIUM WITH WATER)

A

metal + water → metal hydroxide + hydrogen

calcium + water → calcium hydroxide + hydrogen

Observation - Metal disappears along with gas fizzing on surface (Hydrogen). Salt (CaOH) wouldn’t dissolve well –> forms solid white ppt –> makes solution cloudy.

(Some less reactive metals, eg. magnesium, won’t react with water but will react with steam)

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5
Q

Metal reacting with oxygen (eg. sodium + oxygen)

A

metal + oxygen → metal oxide

sodium + oxygen → sodium oxide

(Sodium so reactive it has to be stored in oil to stop it from reacting with oxygen)

Observation - Shiny metal colour on suface becomes more dull (tarnished) over time. White oxide layers form on the surface. The more reactive the element the faster this reaction will be.

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6
Q

Properties of non-metals

A

Don’t conduct heat/electricity, generally low-melting points, brittle, dull

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7
Q

Reactivity series order

A

Most reactive

  1. Potassium
  2. Sodium
  3. Lithium
  4. Calcium
  5. Magnesium
  6. Aluminium
  7. Carbon
  8. Zinc
  9. Iron
  10. Hydrogen
  11. Copper
  12. Silver
  13. Gold

Least reactive

First 4 (K, Na, Li, Ca) react with cold water

5,6,7,8, (Mg, Al, C, Zn) React with steam only

Last 5 (Fe, H, Cu, Ag, Au) don’t react with water

First 10 react with acids
Copper, SIlver and Gold don’t react with acids.

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8
Q

Reactivity series mnemonic

A

Please
Send
Little
Charlie’s
Monkeys
And
Chubby
Zebras
In
Hinged
Cages
Securely
Gaurded

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9
Q

Potassium reacting with cold water

A
  • Lilac flame produced (bc there is enough heat to ignite the hydrogen gas produced)
  • Reaction so vigorous the metal may jum around on the water surface
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10
Q

Sodium reacting with cold water

A
  • Sodium floats on water
  • Effervence can be seen as hydrogen gas is released
  • Moves around fast on water surface bc of amount of hydrogen given off
  • A lot of heat is given off
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11
Q

Magnesium reacting with steam

A
  • Forms hydrogen gas + Magnesium hydroxide (white solid)
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12
Q

Magnesium reacting with dilute acid

A
  • Bubbles would form on the surface and the metal would slowly disappear.
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13
Q

Iron reacting with dilute acid

A
  • Hydrogen gas formed
  • Reaction so slow it would barely be possible to observe or see anything.
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14
Q

Copper, Silver, Gold reacting with dilute acid

A

No reaction

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15
Q

Displacement reaction

A

A reaction where a less reactive element is removed from a compound by a more reactive element.

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16
Q

Displacement reaction with metals

A

When a metal reacts with the ion compound of another metal, a displacement reaction occurs. The more reactive metal reduces the metal ion + makes it gain electrons to form a metal atom again.

The more ractive metal therefore forms a positive ion.

17
Q

Reactivity of Aluminium

A
  • Aluminium = highly reactive but also corrosion resistant
  • Aluminium oxide formed when aluminium is exposed to air
  • Creates super thin layer of oxide (barrier) between metal + air –> whenever metal is scratched –> new layer forms.
  • Aluminium oxide = protective layer which stops corrosion
18
Q

Extracting metals from ores - Two different ways

A

Metals are reactive, and are therefore found in ores (having fomed compounds with other elements).

Hematite - Iron
Bauxite - Aluminium

Two methods of extraction…

  • Displacement reactions in a blast furnace
  • Molten electrolysis

Metals less reactive than carbon, get displaced in a blast furnace.

More reactive metals are extracted through molten electrolysis as it is more difficult to displace the metal using a more reactive element.

19
Q

How do blast furnaces (iron) work?

A

1) Heat the reaction mixture to a very high temperature

2) Provide a chemical which can reduce the iron oxide to iron (aka displace it)

3) Remove other impurities found in the ore, like sand, that would harm the quality of the metal being extracted.

20
Q

Stages of iron refining in the blast furnace…

A
  1. Hematite, coke and limestone is added at the top of the blast furnace
  2. Hot air is blasted into the furnace + combusts with coke to form carbon dioxide.
  3. Carbon dioxide reacts with coke to form carbon monoxide (reducing agent)
  4. Carbon monoxide reacts + reduces Iron (III) oxide to Iron.

REMOVING OTHER IMPURITIES….

  1. Limestone thermally decomposes into more reactive calcium oxide –> reacts with the silicon dioxide in the iron to form slag.
  2. Sulfur impurities in iron react with hot air to form sulfur dioxide.
  3. Coke dissolved in molten iron, reacts with oxygen to form CO2
  4. Molten iron collected from bottom of furnace
  5. Less dense slag sits on top of iron and is also removed (used in road construction)
21
Q

All the reactions within a blast furnace… (word equations)

A

Reducing iron…

carbon + oxygen → carbon dioxide
carbon dioxide + carbon → carbon monoxide
iron(III) oxide + carbon monoxide → iron + carbon dioxide

Removing impurities…

carbon + oxygen → carbon dioxide
sulfur + oxygen → sulfur dioxide
calcium carbonate → calcium oxide + carbon dioxide
calcium oxide + silicon dioxide → slag

22
Q

All the reactions within a blast furnace… (symbol equations)

A

Reducing iron…

C (s)+O² (g)→CO² (g)
S (s) + O² (g) → SO² (g)

Removing impurities…

C (s)+O² (g)→CO² (g)
CaCO³ (s)→CaO (s)+CO² (g)
CaO (s)+SiO² (s)→CaSiO³ (l)

23
Q

How is aluminium extracted from its ore?

A

Aluminium ore (bauxite) contains aluminium oxide, Al2O3, (containing Al³⁺ and O²⁻ ions).

Aluminium can’t be extracted from its ore (Bauxite) using displacement reactions in a blast furnace, because its reactivity is higher than carbon.

→ Instead electrolysis of aluminium oxide (the main component of bauxite) is used

24
Q

Issues with using electrolysis to extract aluminium…

A

Hydrogen = less reactive than aluminium so aqeuous electrolysis, cannot be used.

Melting point of aluminium oxide = too high –> Has to be dissolved in molten cryolite (an ionic compound) for molten electrolysis to occur.

–> Melting point of cryolite/ Al2O3 mixture = lower than melting point of steel (which is the bath in which electrolysis is carried out)

25
Q

What occurs during the molten electrolysis of Al2O3?

A

Cathode - Aluminium

Al³⁺ ion attracted to negative cathode (gain 3 electrons + discharge as aluminium metal)

Al³⁺ + 3e⁻ → Al

Molten aluminium = more dense than cryolite + aluminium oxide solution → Sinks to the bottom + forms a layer of liquid metal which can then be removed)

Anode - Oxygen

Oxygen forms at anodes (gain 2 electrons)

2O²⁻ → O2 + 4e⁻

Because of the high temperatures used, oxygen reacts with the carbon anodes to produce carbon dioxide gas. → causes the anodes to erode away and have to be replaced.

26
Q

Overall reaction of aluminium extraction…

A

2Al²O³ (l)→4Al (l) + 3O² (g)

Chemistry (13 decks)

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