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Chemistry > Module 3.1 - The Periodic Table > Flashcards

Module 3.1 - The Periodic Table Flashcards

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1
Q

Describe the relationship in elements in the same period.

A

-Trends are repeated for each (periodicity)
-Atomic number increases
-

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2
Q

What is the relationship between elements in the same group?

A

Same physical and chemical properties (as same outer electrons + type of orbitals). Repeating pattern of similarity caused by underlying repeating pattern in electron configuration

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3
Q

Put each orbital in order from lowest energy to highest.

A

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4
Q

Put each orbital in order of ascending energy up to 4d.

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d

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5
Q

The 4s orbital has a lower energy than the 3d orbital. What does this mean?

A
  • 4s fills before 3d

- Empties before 3d during ionisation

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6
Q

How can you shorten electron configuration?

A

Use the noble gas before the atom and add any extra orbitals

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7
Q

What is ionisation?

A

When atoms gain or lose electrons to form ions

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8
Q

What is the equation for the first ionisation of sodium?

A

Na(g) –> Na+(g) + e-

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9
Q

What factors affect ionisation energy?

A
  • Atomic radius
  • Nuclear charge
  • Electron shielding or screening
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10
Q

How does atomic radium affect the ionisation energy?

A
  • Larger atomic radius = smaller nuclear attraction to outer electrons
  • Positive nuclear charge is further away from negative electrons
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11
Q

How does nuclear charge affect the ionisation energy?

A

-Higher nuclear charge = stronger attraction

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12
Q

How does the electron shielding affect ionisation energy?

A
  • Inner shells of electrons repel outer electrons as all negative so easier to lose
  • Repelling effect is called electron shielding/screening
  • More inner shells there are, larger the shielding effect + smaller nuclear attraction experiences by outer electrons
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13
Q

What is successive ionisation energy?

A

Measure of the amount of energy required to remove each electron in turn

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14
Q

Why is each successive ionisation energy higher than the one before?

A
  • As each electron is removed, there’s less repulsion between remaining electrons + each shell will be drawn slightly closer to nucleus
  • +ve nuclear charge outweighs -ve charge every time an electron in removed
  • As distance of each electron from nucleus decreases slightly, nuclear attraction increases. More energy required to remove each successive electron
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15
Q

Which element in each period has the highest ionisation energies?

A

-Noble gas as they have a full outer shell of electrons + high positive attraction from nucleus

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16
Q

What trends are there across each period?

A
  • Number of protons in nucleus increases so higher attraction to electrons
  • Atomic radius decreases (as more electrons so more attraction)
  • Shielding is around the same as same inner shells
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17
Q

Why is there a slight decrease in first ionisation energy between groups 2 and 13?

A
  • Group 2’s highest energy electron: s orbital
  • Group 13’s highest energy electron: p orbital
  • P orbitals have a little more energy so are slightly further away from nucleus so electrons are slightly easier to remove
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18
Q

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

A

-Between groups 13 and 15 each electron gets its own p-orbital but from 16 they start pairing in each orbital so slightly easier to remove (as slight repulsion) hence slightly lower first ionisation energy

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19
Q

Why is there a sharp drop in ionisation energy between the noble gas of a period and the group 1 element of the next?

A
  • Increased atomic radius as another shell so further away from nucleus
  • Increase in electron shielding of outermost shell by inner electrons
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20
Q

Why does the first ionisation energy decrease as you go down a group?

A
  • More shells so greater distance between outer electron + nucleus so weaker attraction to outer electrons
  • More inner shells so shielding effect on outer electrons increases so weaker attraction
  • Increase in shielding outweighs the increase in nuclear charge
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21
Q

What are the trends as you go down a group?

A
  • Number of shells increases
  • Shielding increases
  • Atomic radius increases
  • First ionisation energy decreases
22
Q

Describe the structure of metallic lattices.

A
  • Delocalised electrons spread through whole structure
  • Electrons can move within the structure
  • Impossible to tell which electron originated from which cation
  • Charges must balance over the whole structure
  • Cations in fixed positions in the lattice
  • Outer shell electrons are delocalised
23
Q

What are the properties of giant metallic lattices?

A
  • High melting + boiling points
  • Good electrical conductor
  • Malleability and ductility
24
Q

Why do giant metallic lattices have high melting and boiling points?

A
  • Electrons free to move throughout structure but cations remain where they are
  • Attraction between cations + delocalised electrons is very strong
  • High temperature needed to overcome metallic bonds to dislodge cations from rigid positions in lattice
25
Q

Why do metallic lattices conduct electricity?

A
  • Delocalised electrons can move anywhere in the metallic lattice
  • Allows solid to conduct electricity, even when solid
26
Q

Metallic lattices are malleable and ductile. What does this mean and why do they behave this way?

A
  • Ductile: can be drawn out or stretched, permitting metals to be drawn into wires
  • Malleable: can be hammered in different shapes. Metals can be hammered into shapes or hammered in thin sheets
  • Delocalised electron can move giving structure a degree of give, so layers can slide over each other
27
Q

Is silicon considered a metal or a non-metal?

A
  • Si has shiny appearance of metal but is very brittle
  • Si conducts electricity but very poorly
  • ‘In between’ element, usually classified as a metalloid or semi-metal
28
Q

Describe the trend in melting point across a period.

A
  • 1-14: increases steadily. If metallic: nuclear charge increases + so does no. of outer electrons so stronger attraction. If giant covalent: more electrons in covalent bond as go across
  • 14-15: sharp decrease in melting point as simple covalent so weak London forces
  • 15-18: melting remains low due to simple structure
29
Q

As you go across the period of elements forming metallic lattices, what are the trends?

A
  • Ionic charge increases
  • Ionic size decreases
  • No. of outer shell electrons increases
  • Attraction increases: melting + boiling point increases
30
Q

What are the physical properties of group 2 elements?

A
  • Reasonably high melting + boiling points
  • Light with low densities
  • Form colourless (white) compounds
31
Q

Describe generally how group 2 elements react with other elements.

A
  • They’re reactive metals
  • Strong reducing agents
  • Oxidised in their reactions to form 2+ ions
  • M –> M+ + e-
  • M+ –> M2+ + e-
32
Q

Describe the reaction between group 2 elements and oxygen.

A

-React vigorously with oxygen
-Redox reaction
-Product: ionic oxygen with the formula MO
-e.g. calcium + oxygen –> calcium oxide
2Ca(g) + O2(g) –> 2CaO(s)
-calcium oxidised as oxidation number increased from 0 to +2
-oxygen reduced as oxidation number decreased from 0 to -2

33
Q

Describe the reaction between group 2 elements and water.

A

-All (except beryllium) react to form hydroxide M(OH)2 and H2(g)
-Mg reacts slowly with water. Moving down group = more vigorous reaction
-e.g. calcium + water –> calcium oxide + hydrogen
Ca(s) + 2H2O(l) –> Ca(OH)2(aq) + H2(g)
-Calcium oxidised as oxidation number increased from 0 to +2
-Hydrogen reduced as oxidation number decreased from +1 to 0
-Redox reaction

34
Q

Describe the reaction between group 2 elements and dilute acids.

A

-All (except Be) react to form a salt + hydrogen
-Becomes more vigorous as you go down the group
-e.g. calcium + hydrochloric acid –> calcium chloride + hydrogen
Ca(s) + 2HCl(aq) –> CaCl2 + H2
-Calcium oxidised as oxidation number increases from 0 to +2
-Hydrogen reduced as oxidation number decreases from +1 to 0
-Redox

35
Q

Describe the reactions between group 2 oxides and water.

A
  • React with water to form metal hydroxides
  • General equation: MO(s) + H2O(l) –> M(OH)2(aq)
  • Metal hydroxides soluble in water forming alkaline solutions with water as they release OH- ions
  • Typical ph: 10-12
36
Q

Describe the solubility of group 2 metal hydroxides.

A
  • Increases down the group. More soluble hydroxide = release more OH- ions making more alkaline solution so higher pH
  • Be: BeO is insoluble in water
  • Mg: forms Mg(OH)2(s), only slightly soluble in water, dilute, comparatively low OH-(aq) concentration
  • Ba(OH)2(s) more soluble so higher OH-(aq) conc so more alkaline
37
Q

What are some uses of group 2 compounds?

A
  • Neutralising acidic soils: Ca(OH)2 used by farmers + gardeners as lime to reduce acidity of soil
  • Indigestion remedies: indigestion is build up of too much HCl. E.g. milk of magnesia contains Mg(OH)2 neutralising excess acids Mg(OH)2 + HCl –> MgCl2 + 2H2O
  • Building + constructing: CaCO3 useful in building material (in limestone + marble). Used to manufacture glass + steel. Bad as it readily reacts with acids e.g. CaCO3(s) + HCl(l) –> CaCl2(aq) + H2O(l) + CO2(g). Most rainwater has acidic pH so gradual erosion of limestone or marble
38
Q

What are the properties of the halogens?

A
  • Low melting + boiling point

- Exists as diatomic molecules

39
Q

Explain the trend in boiling points down group 17 (the halogens).

A
  • Boiling point increases as you go down the group + physical state changes from gas to liquid to solid
  • Due to successive elements having an extra shell leading to higher level of London forces (more electrons = larger instantaneous dipole so stronger London dispersion forces)
40
Q

Explain the reactivity of the halogens.

A

-Halogens are very reactive + highly electronegative so good oxidising agents (attracting electrons)
-Form halide 1- ions
-As you go down reactivity/oxidising power decreases as:
> Atomic radius increases
> More shielding
> Ability to gain an electron in p sub-shell to form 1- ions decreases

41
Q

Describe and explain the redox reactions of halogens.

A
  • More reactive halogens oxidise + displace halides of a less reactive halogen: a displacement reaction
  • Halogens change colour to indicate a redox reaction has taken place
  • Mixture usually shaken with an organic solvent i.e. cylohexane to distinguish between iodine + bromine
  • Cl2; in water: pale green, in cylohexane: pale green
  • Br2; in water: orange, in cylohexane: orange
  • I2; in water: brown, in cylohexane: violet
42
Q

Describe the displacement reaction between chlorine and bromide ions.

A
  • Chlorine more reactive so oxidises Br-
  • Cl2(aq) + 2Br-(aq) –> 2Cl-(aq) + Br2(aq)
  • Br2 orange in water + cylohexane
  • Chlorine reduced (0 to -1)
  • Bromine oxidised (-1 to 0)
43
Q

Describe the displacement reaction between chlorine and iodide ions.

A
  • Chlorine more reactive so oxidises I-
  • Cl2(aq) + 2I-(aq) –> 2Cl-(aq) + I2(aq)
  • I2 brown in water + violet in cylohexane
  • Chlorine reduced (0 to -1)
  • Iodine oxidised (-1 to 0)
44
Q

Describe the displacement reaction between bromine and iodide ions.

A

-Bromine more reactive so oxidises I-
-Br2(aq) + 2I-(aq) –> 2Br-(aq) + I2(aq)
-I2 brown in water + violet in cylohexane
-Bromine reduced (0 to -1)
Iodine oxidised (-1 to 0)

45
Q

What is a disproportionation reaction?

A

A reaction in which the same element is both oxidised and reduced

46
Q

Describe the reaction of chlorine with water and what this is used for.

A

-Chlorine kills bacteria making water safe to drink
-Reacts with water forming HCl + chloric(I) acid HClO
Cl2(aq) + H2O(l) –> HClO(aq) + HCl(aq)
-Chlorine reduced (0 to -1 in HCl)
-Chlorine oxidised (0 to +1 in HClO)
-Disproportionation reaction

47
Q

Describe the reaction between chlorine and dilute aqueous sodium hydroxide.

A
  • Chlorine only slightly soluble in water + has mild bleaching action
  • Household bleach formed from chlorine + dilute aq NaOH reacting at room temp
  • Disproportionation reaction
  • Cl2(aq) + 2NaOH –> NaCl(aq) + NaClO(aq) + H2O(l)
  • Chlorine reduced (0 to -1 in NaCl)
  • Chlorine oxidised (0 to +1 in NaClO)
48
Q

Describe how household bleach is formed.

A
  • Chlorine only slightly soluble in water + has mild bleaching action
  • Household bleach formed from chlorine + dilute aq NaOH reacting at room temp
  • Disproportionation reaction
  • Cl2(aq) + 2NaOH –> NaCl(aq) + NaClO(aq) + H2O(l)
  • Chlorine reduced (0 to -1 in NaCl)
  • Chlorine oxidised (0 to +1 in NaClO)
49
Q

How do you test for carbonate ions (CO3 2-)?

A
  • Add a dilute strong acid to suspected carbonate
  • Collect any gas for me + pass through limewater
  • Fizzing/colourless gas is produced that turns limewater cloudy (if positive result)
  • CO3 2-(aq) + 2H+(aq) –> H2O(aq) + CO2(g)
50
Q

How do you test for sulfate ions (SO4 2-)?

A
  • Add dilute HCl and BaCl2 to suspected sulfate
  • White precipitate of barium sulfate is produced (if positive result)
  • Ba2+(aq) + SO4 2-(aq) –> BaSO4(aq)
51
Q

How do you test for halides?

A
  • Dissolve suspected halide in water
  • Add an aqueous solution of silver nitrate
  • Note colour of precipitate formed
  • If colour is hard to distinguish add aqueous ammonia (first dilute then concentrated)
  • Note solubility of precipitate in aqueous ammonia
  • Positive test; silver chlorine: white ppt, soluble in dilute ammonia; silver bromide: cream ppt, soluble in concentrated ammonia only; silver iodide: yellow ppt, insoluble in dilute + concentrated ammonia
  • Ag+(aq) + X-(aq) –> AgX(s)
52
Q

How do you test for ammonium ions (NH4+)?

A
  • Add NaOH solution to suspected ammonium compound + warm v gently
  • Test any gas evolved with red litmus paper
  • Ammonia gas turns red litmus paper blue + has a distinct smell (ammonia gas hazardous - done with care)
  • NH4+(aq) + OH-(aq) –> NH3(aq) + H2O(aq)

Chemistry (13 decks)

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