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13A/Ch1 > 3.1 The periodic Table > Flashcards

3.1 The periodic Table Flashcards

1
Q

How are the elements arranged in a periodic table?

A

They are arranged in the order of increasing atomic numbers

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2
Q

What is a period on a periodic table?

A

The horizontal rows in the periodic table

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3
Q

What is a group on a periodic table?

A

The vertical columns

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4
Q

Who’s table is the modern one based off

A

Dmitri Mendeleev

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5
Q

Why did Mendeleev leave gaps in his table?

A

left gaps where no element at the time fitted the repeating patterns. properties of these were predicted

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6
Q

What is meant by periodicity?

A

The repeating trends in chemical and physical properties

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7
Q

What change happens across each period?

A

Elements change from metals to non metals

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8
Q

How can the electron configuration be written in short?

A

The noble gas before (period above)the element is used to abbreviate
E.g Li → 1s22s1 ; Li → [He] 2s1

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9
Q

What are the S block elements?

A

groups 1&2

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10
Q

What are the D block elements?

A

transition metals(middle part)

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11
Q

What are the P block elements?

A

groups 5-8

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12
Q

Define first ionisation energy

A

The energy required to remove 1 mole electron from 1 mol of gaseous atom to form one mole of gaseous 1+ ions

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13
Q

Write an equation for the first ionisation energy

A

X(g)→X+ (g) + e-

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14
Q

What are the factors that affect ionisation energy?

A
  • Atomic radius(further away=less attraction)
  • Nuclear charge
  • Electron shielding or screening (more shells= more shielding)
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15
Q

How does electron shielding work?

A

Inner electron shells repel the outer electron shells. Shielding ‘blocks’ nuclear attraction reaching outer shells

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16
Q

What is the general trend for 1st ionisation across a period?

A

increase as

• across period protons increase but not shells so more nuclear attraction but similar shielding

17
Q

Why does first ionisation energy decrease between group 2 to 3?

A
  • group 2 outer electrons in s orbital,

* group 3 in P orbital so electrons are easier to remove as slightly further away

18
Q

Why does first ionisation energy decrease between group 5 to 6?

A

group 5 electrons in P orbital are single electrons group 6 are spin paired, with some repulsion. so they are slightly easier to remove

19
Q

Does first ionisation increase or decrease between the end of one period and the start of next? Why

A

Decrease
• There is increase in atomic radius
• Increase in electron shielding (new shell added)

20
Q

Does first ionisation increase or decrease down a group?

Why?

A

Decrease
• Shielding increases → weaker attraction
• Atomic radius increases → distance between the outer electrons and nucleus increases → weaker attraction
• Increase in number of protons is outweighed by increase in distance and shielding

21
Q

What are the properties of giant metallic lattices? (4)

A
  • High melting and boiling point
  • Good electrical conductors
  • Malleability(can be moulded)
  • Ductility (can be stretched)
22
Q

Describe the structure, forces and bonding in every element across period 2

A
  • Li & Be → giant metallic ; strong attraction between cations and delocalised electrons ; metallic bonding
  • B & C → giant covalent ; strong forces between atoms covalent
  • N2 ,O2 ,F2 ,Ne → simple molecular ; weak intermolecular forces, covalent bonding
23
Q

Describe the structure, forces and bonding in every element across period 3

A
  • Na, Mg, Al → giant metallic lattice, strong attraction between cations and delocalised electrons, metallic bonding
  • Si → giant covalent lattice, strong forces between atoms, covalent bonding
  • P4 , S8 , Cl2 , Ar → simple molecular, weak intermolecular forces ,covalent bonding
24
Q

Group 2 Properties

A

• high-ish melting & Boiling point
• increasing atomic radii down group
• increasing reactivity with water down group
Increasing solubility
• 2+ oxidation states in compounds (outer shell 2 electrons

25
Q

What happens to group 2 ionisation energies down the group?

A

1st ionisation energy decreases down group

26
Q

What is the reaction with group 2 oxides and water?

A

MO + H20 → M(OH2)
forms water soluble metal hydroxides that form alkaline solutions
lower down group = more Alkaline

27
Q

What are some uses of group 2 Metal compounds?

A
  • neutralise acidic soil ; Ca(OH2)
  • indigestion remedies ; Mg(OH2)
  • metal carbonates for buildings
28
Q

Group 7 Properties

A
  • Low melting & boiling
  • boiling point (solidness) increases down group
  • Reactivity increase UP the group
  • very reactive & electronegative
  • Diatomic
29
Q

Halogen colour in water?

A
  • Cl2; pale green in water
  • Br2; orange in water
  • I2; brown in water (violet in cyclohexane)
30
Q

What’s the order for ion testing, why?

A
  1. carbonate
  2. sulfate
  3. Halide
    otherwise might mistake result for other test (both carbonate and sulfate form precipitates)
31
Q

How do you carry out the carbonate test?

A
  1. Add strong acid to suspected carbonate
  2. Pass gas produced through limewater
  3. If gas produced and water turns cloudy then positive test
32
Q

How do you carry out the Sulfate test?

A
  1. Add dilute HCl & BaCl to suspected sulfate
    2.if sulfate ions present then white precipitate form
    (barium sulfate salt forms)
33
Q

How do you carry out the Halide test?

A
  1. Dissolve halide in water
  2. Add aqueous silver nitrate
  3. Note colour of precipitate formed
34
Q

What colours do the Halide test make?

A

AgCl-white precipitate- Nh3 soluble
AgBr-cream precipitate- strong NH3 only
AgI- yellow precipitate- Nh3 Insoluble

35
Q

How do you carry out the Ammonium test?

A
  1. Add sodium hydroxide to suspected ammonium
  2. heat mixture gently
  3. test gas over ret litmus, turns it blue
    ammonia also has certain smell
36
Q

Uncertainty eq?

A

(absolute uncertainty *2)/ value measured

13A/Ch1 (6 decks)

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