Module 3.1 - The Periodic Table Flashcards
(52 cards)
1
Q
Describe the relationship in elements in the same period.
A
-Trends are repeated for each (periodicity)
-Atomic number increases
-
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
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
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
6
Q
How can you shorten electron configuration?
A
Use the noble gas before the atom and add any extra orbitals
7
Q
What is ionisation?
A
When atoms gain or lose electrons to form ions
8
Q
What is the equation for the first ionisation of sodium?
A
Na(g) –> Na+(g) + e-
9
Q
What factors affect ionisation energy?
A
- Atomic radius
- Nuclear charge
- Electron shielding or screening
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
11
Q
How does nuclear charge affect the ionisation energy?
A
-Higher nuclear charge = stronger attraction
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
13
Q
What is successive ionisation energy?
A
Measure of the amount of energy required to remove each electron in turn
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
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
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
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
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
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
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
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
Why do metallic lattices conduct electricity?
- Delocalised electrons can move anywhere in the metallic lattice
- Allows solid to conduct electricity, even when solid
26
Metallic lattices are malleable and ductile. What does this mean and why do they behave this way?
- 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
Is silicon considered a metal or a non-metal?
- 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
Describe the trend in melting point across a period.
- 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
As you go across the period of elements forming metallic lattices, what are the trends?
- Ionic charge increases
- Ionic size decreases
- No. of outer shell electrons increases
- Attraction increases: melting + boiling point increases
30
What are the physical properties of group 2 elements?
- Reasonably high melting + boiling points
- Light with low densities
- Form colourless (white) compounds
31
Describe generally how group 2 elements react with other elements.
- They're reactive metals
- Strong reducing agents
- Oxidised in their reactions to form 2+ ions
- M --> M+ + e-
- M+ --> M2+ + e-
32
Describe the reaction between group 2 elements and oxygen.
-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
Describe the reaction between group 2 elements and water.
-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
Describe the reaction between group 2 elements and dilute acids.
-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
Describe the reactions between group 2 oxides and water.
- 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
Describe the solubility of group 2 metal hydroxides.
- 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
What are some uses of group 2 compounds?
- 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
What are the properties of the halogens?
- Low melting + boiling point
| - Exists as diatomic molecules
39
Explain the trend in boiling points down group 17 (the halogens).
- 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
Explain the reactivity of the halogens.
-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
Describe and explain the redox reactions of halogens.
- 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
Describe the displacement reaction between chlorine and bromide ions.
- 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
Describe the displacement reaction between chlorine and iodide ions.
- 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
Describe the displacement reaction between bromine and iodide ions.
-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
What is a disproportionation reaction?
A reaction in which the same element is both oxidised and reduced
46
Describe the reaction of chlorine with water and what this is used for.
-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
Describe the reaction between chlorine and dilute aqueous sodium hydroxide.
- 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
Describe how household bleach is formed.
- 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
How do you test for carbonate ions (CO3 2-)?
- 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
How do you test for sulfate ions (SO4 2-)?
- 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
How do you test for halides?
- 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
How do you test for ammonium ions (NH4+)?
- 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)
