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chemistry > periodic table and energy 3.1 > Flashcards

periodic table and energy 3.1 Flashcards

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1
Q

\what 3 factor does nuclear attraction depend on?

A
  • atomic radius
    -nuclear charge
    -electron shielding or screening
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2
Q

describe the history of the periodic table?

A
  • early 1800s they ordered the periodic table according to atomic mass
    -1817- dobereiner’s triads, certain group of 3 elements ordered by atomic weight would have a middle element with a weight and properties that were roughly an average of the other two elements.
    -1863- Newland octaves. noticed every 8th element ordered by mass had similar properties. broke down when some transition metals didn’t fit his pattern,
    -1869= Mendeleev, left gaps where elements didn’t fit
    -modern periodic table- order by proton number
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3
Q

first ionisation energy

A

the energy needed to remove one mole of electrons from one mole of atoms in their gaseous state to form one mole of 1+ ions(also in their gaseous state)

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4
Q

the effect of shielding on ionisation energies

A
  • the more electron shells between the positive nucleus and the negative outer electrons the weaker the attraction. therefore less energy is required to remove an electron. (ionisation decrease)
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5
Q

the effect of nuclear charge on ionisation energies.

A
  • higher nuclear charge means that there are more protons in the nucleus, so there’s a bigger attraction between the nucleus and the outer electron. Therefore more energy is required to remove the electron(ionisation energies increase)
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6
Q

trend in ionisation energy down a group

A

-decreases because the outer electrons are further away from the nucleus so force of attraction are weaker so the energy required to remove one electron decreases. shielding also increases and has the same effect.

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7
Q

trend in ionisation energy across a period

A

-it increases as the number of protons in the nucleus increases so there is a higher nuclear attraction. shielding does not affect the ionisation energy as its the same.

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8
Q

why is the first ionisation energy of aluminium less than that of magnesium, even though aluminium has a higher nuclear charge?

A

-magnesium sits in the s orbital whereas aluminium sits in the p orbital which is a higher energy subshell so less energy is needed to remove an electron.

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9
Q

why does sulphur have a lower ionisation energy than phosphorus ?

A

sulphur electrons are paired in there 3p orbital, while phosphorus is singularly occupied in the orbital. Mutual repulsion between paired electrons means less energy is required to remove an electron.

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10
Q

metallic bonding

A

strong electrostatic attraction between cations (positive ions) and delocalised electrons.

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11
Q

properties of giant metallic structures

A
  • high melting and boiling points due to the high electrostatic forces between positive ions and the negative delocalised electrons
    -good electrical conductor-as the delocalised electrons are free to move and carry change.
    -insoluble in water-as the metallic bond is too strong
    -malleable- due to the delocalised electrons can move around
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12
Q

why does magnesium have a higher melting point than sodium?

A
  • it can donate 2 electrons whereas as sodium can only donate one. Magnesium ions have a stronger positive charge so the attraction to the delocalised electrons is stronger. There are also more delocalised electrons in magnesium than there is in sodium.
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13
Q

structure and properties of graphite

A
  • giant covalent structure, each carbon bonded 3 times with the 4th electron delocalised,
    -layer can slide easily due to weak forces between thee layers
    -can conduct electricity due to delocalised electrons
  • high melting points-strong covalent bonds+insoluable
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14
Q

structure and properties of diamond and silicon

A

-giant covalent structure
- bonded to 4 carbons in a tetrahedral shape.
-high melting and boiling point due to strong covalent bonds
-doesn’t conduct electricity as it has no delocalised electrons
-insoluble

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15
Q

structure and properties of graphene

A
  • 1 layer of graphite,1 atom thick in a hexagon carbon ring
    -lightweight and transparent as its 1 atom thick
    -excellent conductor of electricity due to delocalised electrons
    -high strength property due to strong covalent bonds
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16
Q

two uses of graphene

A

-highspeed computers
-smart phone screens

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17
Q

variations in melting points across period 2 and 3 in terms of structure and bonding

A
  • melting point begins to increase, the 3 elements are metals except boron and have a giant metallic structure so have strong forces between positive and delocalised electrons. nuclear charge also increases
  • carbon and silicon are the peak in melting points as they have giant covalent structures so have strong forces
    -last 4 elements have low melting points as they are have a simple molecular structure so have weak induced dipole-dipole force between molecules
    -phosphorous causes a peak as it is p4 which has a higher so stronger induced dipole-dipole forces.
  • Ar is a single atom so melting point is low
18
Q

properties of group 2 elements

A
  • high melting and boiling points
    -light metals with low densities
    -form colourless(white) compounds
19
Q

what ions do group 2 elements form in a redox reaction?

A

2+ ions

20
Q

trend in reactivity down group 2

A

reactivity increases down group two due to more shielding and the atomic radius increases therefore there is a weaker attraction between the nucleus and outer electrons so an electron is lost more easily

ionisation energy decreases therefore

21
Q

group 2 elements increases in positive charge as we go down, why then is the first ionisation energy decreasing?

A

the shielding effect overrides an increases in positive charge

22
Q

how do group 2 elements react with water?

A
  • they react vigorously with water forming a metal hydroxide M(OH)2and hydrogen gas. except from (Be)
  • is a redox reaction
23
Q

how does magnesium react with water?

A

slowly producing magnesium oxide(Mgo) instead of hydroxide
- is a redox reaction

24
Q

how do group 2 elements react with oxygen?

A
  • vigorously forming metal oxides
    -is a redox reaction
25
Q

How do group 2 elements react with dilute acids?

A
  • vigorously to form salt and hydrogen gas
  • redox reaction
26
Q

how do group 2 oxides react with water?

A
  • they form metal hydroxides which are soluble in water, and form alkaline solution with water as they release OH- ions. The typical PH of theses solutions is between 10 and 1
27
Q

trend in solubility and alkalinity of group 2 metal hydroxides

A

solubility of the hydroxides in water increases down group 2. more soluble hydroxides release more OH- ions, and so will make a more alkaline solution, with a higher PH.

28
Q

uses of group 2 compounds

A
  • calcium hydroxide Ca(OH)2 is used to neutralise soil
    -Magnesium hydroxide Mg(OH)2 is used in indigestion remedies to neutralise excess acid in the stomach producing salt and water
  • calcium carbonate CaCO3 used in building and construction such as limestone and marble.
29
Q

what do halogens exist as?

A

-diatomic molecules

30
Q

what is the outer shell configuration of group 7?

A

-s2p5

31
Q

trend in boiling points of group p7

A

the boiling points increases down a group as the molecules have more electrons so have stronger induced dipole-dipole forces. Therefore Cl2 is a gas at room temperature and pressure, Br2 is a liquid and F2 is a solid

32
Q

what is a halide ion?

A

when a group 7 element gains an electron in a redox reaction to form a 1- ion

33
Q

trend in reactivity of group 7 elements

A

-reactivity decreases down the group as group 7 elements react by gaining an electron. Shielding and atomic radius increases down a group meaning that the electron is less attracted to the nucleus.

34
Q

disproportionation reaction

A

oxidation and reduction of the same element

35
Q

uses of chlorine

A

-water purification. Reacts with water forming HCL and HCLO (chloric acid)
-forms bleach when it reacts with dilute aqueous sodium hydroxide. forming Nacl, NaCLO and H20

36
Q

what are the benefits and risks of using chlorine for water purification

A
  • Benefits: it kills bacteria
    -Risks: toxic gas, reacts with hydrocarbons producing chlorinated hydrocarbons which are dangerous
37
Q

how do we test for halide ions?

A
  • add an aqueous solution of silver nitrate to the dissolve suspected halide ion and observe the colour change. as different halide ions react form different colour silver halide precipitates. If the colour is hard to distinguish add aqueous ammonia and note the solubility of the ammonia in the precipitate.
38
Q

colours the halide ions form when added to silver nitrate

A
  • silver chloride white and soluble in dilute ammonia
  • silver bromide is cream and only soluble in concentrated ammonia
    -silver iodide is yellow and insoluble in both dilute and concentrated ammonia
39
Q

How do we test for for carbonate ions ( Co3 2-)

A

Co3 2- ions react with acids so to test for them…

Add a dilute strong acid to the suspected carbonate and collect any gas formed and pass it through limewater. If the Co3 2- ion is present fizzing/colourless gas is produced which turns limewater cloudy.

40
Q

How do we test for sulfate( so4 2-) ions?

A
  • sulfate ions react with barium ions to form barium sulfate( BaSO4) so test for sulfate ions we…..

Add dilute HCL and barium chloride to suspected sulfate and if present a white precipitate of barium sulfate is produced.

41
Q

How do we test ammonium ions (NH 4+)

A
  • ammonium ions react with hydroxide ions to produce ammonia and water. To test for them:
  • add sodium hydroxide solution to the suspected ammonium compound and warm very gently. Test any gas present. If NH 4 + is present it will turn red litmus paper blue it us also has a distinctive smell

chemistry (9 decks)

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