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Chemistry: Module 3 > 3.1 The periodic table > Flashcards

3.1 The periodic table Flashcards

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1
Q

How are elements in the periodic table arranged?

A
  • by increasing atomic
    (proton) number
  • in periods showing repeating trends in physical and chemical properties
  • in groups having similar chemical properies
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2
Q

What is the periodic trend in electron configuration across periods 2 and 3?

A

Across period 2 - 2 electrons fill the 2s sub-shell before 6 electrons fill the 2p sub-shell Across period 3 - 2 electrons fill the 3s sub-shell before 6 electrons fill the 3p sub-shell

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3
Q

Definition of ‘first ionisation energy’

A

the energy required to remove 1 electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions

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4
Q

What are the factors affecting ionisation energy

A

atomic radius
nuclear charge
electron shielding
nuclear attraction

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5
Q

What are the trends of first ionisation trends across a period

A

atomic radius slightly decreases
electron shielding stays the same
nuclear charge increases as there are more protons nuclear attraction increases across so a higher ionisation energy needed

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6
Q

What are the trends of first ionisation trends down a group?

A

atomic radius increases - more shells
electron shielding increases - more shells
nuclear charge decreases
weak nuclear attraction so less energy required to remove an electron
ionisation energy decreases down a group

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7
Q

Describe the variation of melting points across periods 2 and 3?

A

groups 1 - 14 melting points increase steadily, giant metallic (higher nuclear charge, strong attraction) and covalent structure

group 14 - 15 sharp decrease in melting points, simple molecular structures - weak intermolecular structures

group 15 - 18 melting points stay relatively low, simple molecular structures

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8
Q

Describe variance in first ionisation energy between groups 2 and 3(13)?

A

overall ionisation energy increases across a period between groups 2 and 13 (Be to B, Mg to Al) - slight decrease
group 13 - p orbital
group 2 - s orbital
p orbital have higher energy than s orbitals so they are marginally further away from the nucleus
meaning that electrons p orbitals are slightly easier to remove so elements have lower ionisation energies

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9
Q

Why is there decrease in first ionisation energy between groups 15 and 16?

A

similar decrease to group 2 and 13
15 and 16 - (N to O, P to S)
group 13 - 18 all electrons in p orbital
groups 13, 14, 15 - each p orbital contains a single electron
group 16 - outermost electron is spin paired so there is some repulsion - making electron easier to remove so less ionisation energy needed

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10
Q

What is metallic bonding?

A

strong electrostatic attraction between cations(+) and delocalised electrons

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11
Q

What are metallic bonding properties?

A

high melting/boiling points
ductile
malleable

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12
Q

Define ‘giant metallic lattice’

A

a 3d structure of positive ions and delocalised electrons, bounded by strong metallic bonds

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13
Q

Define ‘giant covalent lattice’

A

a 3d structure of atoms that are all bonded together by strong covalent bonds

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14
Q

What are the structure and properties of diamond?

A

4 strong covalent bonds per atom
tetrahedral shape - 109.4
high melting point
hard
non- conducting

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15
Q

What are the structure and properties of graphene?

A

2d giant lattice
1 carbon atom thick of interlocking hexagonal carbon rings
extremely strong
very light
conducts electricity

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16
Q

What are the structure and properties of graphite?

A

3 strong covalent bonds + 1 delocalised electron
trigonal planar (planes of atoms in hexagons) - 120 high melting point
soft - layers slide
conducts electricity - delocalised electrons

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17
Q

What are the physical properties of giant covalent lattices?

A

high melting and boiling point
non-conductors of electricity - no free charged particles (except graphite)
insoluble in polar and non-polar solvents - covalent bonds too strong

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18
Q

What are the physical properties of giant metallic lattices?

A

high melting and boiling point - strong attraction between positive ions and negative delocalised electrons
good electrical conductivity - delocalised electrons can move freely, even in solid state
malleable and ductile - delocalised electrons can move so the atoms or layers can slide past each other

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19
Q

What are the trends in melting point across period 2 and 3?

A

melting point increases from Group 1 to Group 14 - elements have giant structures (metallic then covalent)
melting points decrease from Group 14 to Group 15 because the structure changes to simple molecular - only held together by weak intermolecular forces

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20
Q

Group 2 metals + oxygen

A

metal oxide
vigorous reaction

21
Q

Group 2 metals + water

A

metal hydroxide + hydrogen gas
beryllium doesnt react
mg reacts very slowly - vigour increasing as you go down
redox

22
Q

Group 2 metals + dilute acid

A

salt + hydrogen gas
beryllium doesnt react
reaction more vigorous going down

23
Q

Where are the s-, p-, d- and f- blocks on the periodic table?

A

s d p
f

24
Q

What are the trends in reactivity down group 2 in terms of 1st and 2nd ionisation energies?

A

2 outer electrons lost more easily going down the group reactivity increases going down the group ionisation energy decreases

25
Q

What is the trend in the increasing alkalinity of group 2 metals?

A

alkalinity increases down group 2
when reacting with water they form more soluble metal hydroxides - releases more OH- ions, making it more alkaline with higher pH

26
Q

How do group 2 metals neutralise acidic soils?

A

calcium hydroxide (‘lime’) - Ca(OH) 2
-reduces acidity level of soil

27
Q

How do group 2 metals help indigestion?

A

in the stomach too much hcl - indigestion
magnesium hydroxide - Mg(OH)2 and calcium carbonate - CaCO3 - acts as an antacid
-neutralises the excess acid and produces salt and water
acid + metal — salt + water

28
Q

What type of molecules are halogens?

A

diatomic

29
Q

Describe the trend in boiling point in halogens?

A

down the group boiling point increases and physical state changes from gas - liquid - solid

because each element has extra shell of electrons - leading to higher level of London forces (induced dipole - dipole) between molecules

more energy needed to remove an electron so increased boiling point
reactivity decreases

30
Q

Describe the electron configuration of halogens/group 7/group 17 ?

A

7 outer shell electrons
xs2 xp5
gain 1 electron to form 1- ions - halide ions

31
Q

Describe the reactivity of halogens?

A

very reactive and highly electronegative - good at attracting electrons
reactivity and oxidising power decrease down the group

atomic radius decreases
electron shielding increases
ability to gain electron from p-subshell to form 1- ion decreases

32
Q

What colour do the 3 halogens change in water?

A

Cl - pale green
Br - orange
I - brown

33
Q

What colour do the 3 halogens change in cyclohexane?

A

Cl - pale green
Br - orange
I - violet

34
Q

Define ‘disproportionation’

A

oxidation and reduction of the same element

35
Q

What is the full equation of a reaction of chlorine with water (disproportionation)?

A

Cl2(aq) + H2O(aq) —> HClO(aq) + HCl(aq)
Chlorine is oxidised in HClO from 0 to +1
Chlorine is reduced in HCl from 0 to -1

36
Q

What is the full equation of a reaction of chlorine with cold, dilute, aqueous NaOH (disproportionation)?

A

Cl2(aq) + 2NaOH(aq) —> NaCl(aq) + NaClO(aq) + H2O(l)
Chlorine is oxidised in NaClO from 0 to +1
Chlorine is reduced in NaCl from 0 to -1

NaClO = bleach

37
Q

What are the benefits of chlorine in water treatment?

A

kills bacteria
eradicated water-carried diseases e.g. cholera and dysentery

38
Q

What are the risks and ethical considerations of using chlorine in water treatment?

A

chlorine is a toxic gas
chlorine can form chlorinated hydrocarbons, causing cancer

peru stopped water treatment - cholera outbreak
people cant choose if they want chlorine in their water or not - unfair

39
Q

What is the ionic equation of chlorine oxidising bromine and the colour change in water and cyclohexane?

A

Cl2(aq) + 2Br-(aq) –> 2Cl-(aq) + Br2(aq)
water: orange cyclohexane: orange

40
Q

What is the ionic equation of chlorine oxidising iodine and the colour change in water and cyclohexane?

A

Cl2(aq) + 2I-(aq) –> 2Cl-(aq) + I2(aq)
water: brown cyclohexane: violet

41
Q

What is the ionic equation of bromine oxidising iodine and the colour change in water and cyclohexane?

A

Br2(aq) + 2I-(aq) –> 2Br-(aq) + I2(aq)
water: brown cyclohexane: violet

42
Q

What is the ionic equation for Cl- with AgNO3 and NH3?

A

Ag+ (aq) + Cl- (aq) —> AgCl(s)

43
Q

What is the ionic equation for Br- with AgNO3 and NH3?

A

Ag+ (aq) + Br- (aq) —> AgBr(s)

44
Q

What is the ionic equation for I- with AgNO3 and NH3?

A

Ag+ (aq) + l- (aq) —> Agl(s)

45
Q

Describe the test for halide ions

A
  1. dissolve halide in water
  2. add silver nitrate (aq)
  3. colour change - Cl(white), Br(cream), I(yellow)
  4. if colour hard to tell - add dilute ammonia
  5. Cl(soluble in dilute), Br(soluble in concentrated), I(insoluble)
46
Q

Describe the test for carbonate ions and give the equation

A

CO3 2-(aq) + 2H+(aq) –> H20(aq)
+ CO2(g)
1. add dilute strong acid to solution
2. pass any gas formed through limewater
3. gas should turn limewater cloudy

47
Q

Describe the test for sulfate ions and give the equation

A

Ba 2+(aq) + SO4 2-(aq) –> BaSO4(aq)
1. add dilute HCl and barium chloride
barium sulfate produces a white precipitate

48
Q

Describe the test for ammonium ions and give the equation

A

NH4 +(aq) + OH-(aq) –> NH3 (aq) + H2O (aq)
1. add NaOH to solution and gently warm
2. test any gas with red litmus paper

ammonia gas - red litmus paper - blue
- has distinctive smell (hazardous)

Chemistry: Module 3 (2 decks)

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