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Remove Chemistry > Topic 2 - Metals and Groups of the Periodic Table > Flashcards

Topic 2 - Metals and Groups of the Periodic Table Flashcards

1
Q

What elements exist as diatomic molecules?

A

Hydrogen (H) , Nitrogen (N), Oxygen (O), Fluorine (F), Chlorine (Cl), Iodine (I), and Bromine (Br)

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2
Q

Why do elements in the same group react similarly?

A

They have the same number of outer electrons, and therefore need to gain/lose the same amount, and so react similarly.

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3
Q
  1. What are the group 1 elements also known as?
  2. Why are they known as this?
A
  1. The Alkali metals
  2. They react with water to form alkaline metal hydroxides
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4
Q

What are the reactions with water of:
1. Lithium (Li)
2. Sodium (Na)
3. Potassium (K)
How alkaline are the pH of each of these?

A
  1. Floats on water, fizzes (Alkaline)
  2. Floats on water, melts into a ball, fizzes, moves rapidly on the surface of water (More Alkaline)
  3. Burns with a lilac flame, melts, moves rapidly on the surface of water (Even more Alkaline)

Therefore, alkaline pH increases as you go down group 1

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5
Q
  1. What is produced when the Alkali metals react with water?
  2. Therefore, what is the formula for the reaction?
A
  1. Metal hydroxides and Hydrogen
  2. Metal + Water –> Metal Hydroxide + Hydrogen
    The (Metal) Hydroxide solution formed is Alkaline (This can be shown using universal indicator)

e.g. Sodium + Water –> Sodium Hydroxide + Hydrogen
2Na+2H2O –>2NaHO+H2

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6
Q

How is reaction shown using dot and cross diagrams?

A

First, draw the elements before the reaction, showing the number of electrons in their outer shells. (On the left of the arrow). Use arrows to show the movement of electrons. Then draw the same but after the reaction (After the arrow). Place square brackets around the ions/ reacted elements, and place the charge outside the brackets in the top right. The charge is how many electrons the element needs to gain/lose to get back to normal, indicated using [number]+/-

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7
Q

Why are elements more reactive as you move down group 1?

A

The outer electrons are further from the nucleus, meaning they have less force of attraction towards the nucleus. Therefore they are easier to lose and the element will react easier.

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8
Q

How are elements arranged in the periodic table?

A

By Atomic Number (Low to High)

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9
Q
  1. How do you write electron configurations?
  2. How do you know what group elements are in based on their electron configuration?
A
  1. You write how many electrons are in each shell, e.g. Magnesium would be 2,8,2 (Total number of electrons = atomic number)
  2. The number of electrons in the outer shell is what group an element is in, e.g. 2,8,2 would be in group 2, and 2,8,7 would be in group 7
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10
Q
  1. Why are the elements in group 0 unreactive?
  2. Why are elements in group 1 so reactive?
A
  1. They already have a full outer shell, so don’t need to react to gain/lose electrons and get a full outer shell
  2. They only have one electron in their outer shell, which is easy to lose
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11
Q

What are elements in group 7 known as?

A

The halogens

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12
Q

What are the state at room temp and colour of:
1. Fluorine (F)
2. Chlorine (Cl)
3. Bromine (Br)
4. Iodine (I)

A
  1. Gas, Light Yellow/Green
  2. Gas, Green
  3. Liquid (which easily vaporises) Red/Orange liquid and gas
  4. Solid (Which easily forms vapour) Grey solid - Forms Purple gas
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13
Q
  1. What are 5 physical properties all halogens have?
  2. What is the rule for melting and boiling point as you go down the group
A
  1. They are all diatomic, have strong covalent bonds, have weak intermolecular forces, low melting + boiling points, and are all poor conductors of heat and electricity.
  2. Melting and boiling point increase as you go down group 7
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14
Q
  1. What is the trend of reactivity as you go down group 7?
  2. Why is this?
A
  1. Reactivity decreases as you go down group 7
  2. The outermost electrons are further from the nucleus as you go down the group, therefore it is harder to attract a new electron (as group 7 elements need to gain 1 electron), as it will be less strongly attracted to the nucleus.
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15
Q

What is the formula for the reactions of:
1. Metal + Oxygen
2. Metal + Water
3. Metal + Acid

A
  1. Metal + Oxygen –> Metal Oxide
    e.g. Sodium + Oxygen –> Sodium Oxide
  2. Metal + Water –> Metal Hydroxide + Hydrogen
    e.g. Sodium + Water –> Sodium Hydroxide + Hydrogen
  3. Metal + Acid –> Salt + Hydrogen
    e.g. Sodium + Hydrochloric Acid –> Sodium Chloride + Hydrogen
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16
Q

What is a displacement reaction?

A

A reaction where a more reactive metal displaces a less reactive metal from a compound.

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17
Q

What are the reactions with cold water and hot acid of:
1. Aluminium (Both)
2. Copper (Both)
3. Iron (Just Cold Water)
4. Magnesium (Just Cold Water)
5. Zinc (Just Cold Water)

  1. Why were some not reacted with acid in this experiment?
A
  1. Cold Water: No Reaction l Hot Acid: Moderate Fizzing
  2. Cold Water: No Reaction l Hot Acid: Slight/No reaction
  3. Cold Water: Slight Fizzing
  4. Cold Water: Moderate FIzzing
  5. Cold Water: Moderate/Slight Water
  6. Some metals were too reactive to be placed in acid, as it would be too dangerous
18
Q

What is the order of the reactivity series? What is an acronym to help?

A

Potassium - Please
Sodium - Stop
Calcium - Calling
Magnesium - Me
Aluminium - A
Carbon - Careless
Zinc - Zebra
Iron - Instead
Tin - Try
Lead - Learning
Hydrogen - How
Copper - Copper
Silver - Saves
Gold - Gold

19
Q

What is Oxidation? What is Reduction?
(Hint: OIL RIG)

A

Oxidation is the Loss of Electrons
Reduction is the Gain of Electrons

20
Q
  1. Electrons on the left of the arrow show they are being…
  2. Electrons on the right of the arrow show they are being…
A
  1. Lost (Oxidised)
  2. Gained (Reduced)
21
Q

What is the half equation (atom to ion) of:
1. Sodium (Group 1)
2. Fluorine (Group 7)
3. Aluminium (Group 3)

A
  1. Na –> Na+ + e-
  2. F2 (diatomic) + 2e- –> 2F-
  3. Al –> Al3+ +3e-
22
Q

In becoming ions, metals always…
In becoming ions, non-metals always…

A
  1. Lose electrons (Oxidation)
  2. Gain electrons (Reduction)
23
Q

What are redox rections?

A

Reactions where both oxidation and reduction occur

24
Q

How do you form redox equations?

A

Put together the two half equations, making sure electrons are equal (Multiplying reactions to make electrons equal if needed), then cancel out the electrons

25
Q

Why does a metal (e.g. Iron) become less capable at carrying current if it comes into contact with oxygen?

A

The metal’s delocalised electrons will react with the oxygen to form metal oxide (e.g. Iron Oxide), meaning there are less delocalised electrons to carry current, and metal oxide is ionic, and so can’t conduct electricity when solid.

26
Q

When you react a Metal with an Acid it forms a Salt and Hydrogen, but what are the specific salts produced for these acids:

  1. Metal + Sulfuric Acid
  2. Metal + Nitric Acid
  3. Metal + Phosphoric Acid
  4. Metal + Hydrochloric Acid
A
  1. Metal Sulfate (SO4^2-)
  2. Metal Nitrate (NO3^-)
  3. Metal Phosphate (PO4^3-)
  4. Metal Chloride (Cl^-)

e.g. Magnesium (Mg) + 2 Hydrochloric Acid (2HCl) –> Magnesium Chloride (MgCl2) + Hydrogen (H2)

27
Q
  1. What is corrosion?
  2. What is rusting?
A
  1. The destruction of materials by chemical reaction with substances in the environment
  2. The name we give to describe the corrosion of iron
28
Q

What does iron react with to rust?

A

Oxygen and Water - BOTH are needed for rusting to occur

29
Q

What is the chemical formula for rust (Word and symbol equation)

A

Hydrated Iron (III) Oxide - The III is needed (Valency of this Iron)
Fe2 O3 . H2O - The . is needed (Shows there is a presence of water)

30
Q

Form the ionic equation from the half equations:
1. Fe –> Fe^3+ + 3e^- (Oxidised)
2. O2 + 4e^- –> 2O^2- (Reduced)

A

Steps:
Fe –> Fe^3+ + 3e^- (x4) = 4Fe –> 4Fe^3+ + 12e^-
O2 + 4e^- –> 2O^2- (x3) = 3O2 + 12e^- –> 6O^2

4Fe –> 4Fe^3+ + 12e^-
3O2 + 12e^- –> 6O^2
PUT TOGETHER

4Fe + 3O2 + 12e^- –> 4Fe^3+ + 12e^- + 2O^2
CANCEL OUT THE ELECTRONS

4Fe + 3O2 –> 4Fe^3 + 2O^2
SIMPLIFY 4Fe^3 + 2O^2 TO Fe2 O3 (Rust formula)

4Fe + 3O2 –> 2 Fe2 O3
FINAL ANSWER

31
Q

What are the three methods for prevention of rust, what do these involve?

A
  1. Barrier - Painting, Oiling, Greasing, Electro Plating. Provides an impermeable barrier between the iron and the air.
  2. Sacrificial Protection - Attach a more reactive metal on top, the more reactive metal oxidises in favour of the iron, so it is a ‘sacrifice’ in order to protect the iron. THe metal needs to be replaced once fully oxidised - before the iron starts to rust
  3. Galvanising - Involves both barrier and sacrificial. The iron/steel is plated in Zinc. The top layer of zinc will oxidise, but no further corrosion will take place. If the zinc gets scratched, a lower layer will oxidise, until the scratch is repaired.
32
Q
  1. How is aluminium extracted from its ore?
  2. What is the product of this reaction?
A
  1. Using electrolysis
  2. Bauxite
33
Q

Why can aluminium not be extracted by heating it’s ore with carbon?

A

Aluminium is more reactive than carbon, so carbon won’t displace aluminium in a displacement reaction.

34
Q
  1. What are native metals?
  2. How are most metals found?
A
  1. A native metal is an unreactive metal that can be found as elements
  2. Most metals are found naturally in rocks called ores. They are in compounds, chemically bonded to other metals.
35
Q

What is the formula (word) for the electrolysis of aluminium oxide?

A

Aluminium Oxide –> Aluminium + Oxygen

36
Q

In a pure metal, what do all the electrons have the same? What does this mean they can do, what is this property called?

A

In a pure metal all atoms are the same shape and size, this means the layers can slide easily, called malleable (can be hammered into shape)

37
Q

How do the atoms differ in an alloy compared to a pure metal? this makes it harder for what? Therefore alloys are … than pure metals.

A

Atoms are different sizes in an alloy, this makes it harder for layers to slide, and so alloys are stronger than pure metals.

38
Q

What is steel an alloy of?

A

Iron + Carbon (Sometimes other metals too)

39
Q

What proportions of what iron is alloyed with in:
Mild Steel
High Carbon Steel
Low Alloy Steel
Stainless Steel

A

Mild Steel - 0.25% Carbon
High Carbon Steel - 0.5-1.4% Carbon
Low Alloy Steel - 1-5% of other metals (Cr,Ni,Ti)
Stainless Steel - 20% Chromium, 10% Nickel

40
Q

What are the uses and most important properties of:
Mild Steel
High Carbon Steel
Low Alloy Steel
Stainless Steel

A

USES:
Mild Steel - Car Body Panels, Wires
High Carbon Steel - Tools + Chisels
Low Alloy Steel - Construction Bridges, High speed tools
Stainless Steel - Cutlery + Sinks, Chemical plants

PROPERTIES:
Mild Steel - Soft and Malleable
High Carbon Steel - Hard
Low Alloy Steel - Hard, Strong, low ductility+malleability
Stainless Steel - Strong + resistant to corrosion

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