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EL Flashcards

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1
Q

What are isotopes?

A

Atoms of the same element with the same atomic number but with a different number of neutrons resulting in a different mass number.

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2
Q

What is the formula for percentage yield?

A

Experimental mass/ theoretical mass x100

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3
Q

What is the formula for percentage composition?

A

Mr of an element/ Mr of compound x100

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4
Q

Formula for concentration

A

Mass or moles/ volume

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5
Q

Formula for concentration

A

Mass or moles/ volume

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6
Q

Volumetric solution simple method for acid base titration

A
  1. Weigh sample bottle of balance
  2. Transfer solid to beaker and reweight sample bottle
  3. Record difference in mass
  4. Add distilled water and stir with glass rod until dissolved
  5. Transfer to volumetric flask
  6. Make up to 250cm3 with distilled water
  7. Shake flask
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7
Q

Acid base titration method

A
  1. Fill burette with standard solution of known conc
  2. Use pipette filler to transfer 25cm3 of solution with unknown conc into conical flask
  3. Add couple drops of indicator
  4. Record initial burette reading
  5. Titrate contents of conical flask by adding solution to it from the burette until the indicator undergoes definite colour change
  6. Record final burette reading and Calc titre volume
  7. Repeat for concordant results
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8
Q

Experimental techniques (titration)

A

Use a balance to measure mass.
Use a volumetric flask to measure volume of a solution.
Use a gas syringe to measure volume of gas.

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9
Q

What are shells?

A

Electrons orbit the nucleus at different energy levels which increase with distance from the nucleus.

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10
Q

What are the seperate sub shells?

A

S (2 electrons), p (8 electrons) d ( 18 electrons) and f (32 electrons)

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11
Q

What are orbitals?

A

Electrons are held in clouds of negative charge.

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12
Q

What is the shape of the s-orbital?

A

Spherical

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13
Q

What is the shape of the p-orbital?

A

Dumbell

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14
Q

What happens in a fusion reaction?

A

Two lighter nuclei join to give a single heavier nucleus. This process releases energy but requires some extreme conditions of high temperature and pressure. This is how some elements are formed.

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15
Q

What is ionic bonding?

A

An electrostatic attraction between positive and negative ions. It occurs between a metal and non- metal. Electrons are transferred from the metal to non- metal to achieve full outer shells. Ionic substances only conduct electricity when molten or aqueous as when they are solid the ions cannot move to carry a charge. When the electrons are transferred it creates charged particles called ions. Oppositely charged ions attract through electrostatic forces to form a giant ionic lattice.

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16
Q

What is a covalent bond?

A

The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.

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17
Q

How do covalent bonds form?

A

Between two non- metals. Electrons are shared between the two outer shells in order to achieve full outer shells. Multiple electron pairs can be shared to form multiple bonds.

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18
Q

Dative bond

A

Both of the electrons in the shared pair are supplied from a single atom. Once a dative bond has formed it is treated as a standard covalent bond because it reacts in the same way. Eg ammonia NH3

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19
Q

What is simple covalent?

A

Substances with a simple molecular structure consisting of covalently bonded molecules help together with intermolecular forces.

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20
Q

How is the shape of a simple molecule or ion determined?

A

By the number of electron pairs around the central atom and the repulsion between them. Each electron pair naturally repels each other so the largest bond angle possible exists between the covalent bonds.

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21
Q

lone pair repulsion

A

Any lone pairs present around the central atom provide additional repulsive forces which changes the bond angle. For every lp present the bond angle reduces by 2.5 degrees.

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22
Q

linear shape

A

2 bonding pairs
180 degrees

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23
Q

bent shape

A

2 bonding pairs and 2 lone pairs
104.5 degrees

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24
Q

trigonal planar shape

A

3 bonding pairs
120 degrees

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25
Q

Trigonal pyramidal

A

3 bonding pairs
1 lone pair
107 degrees

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26
Q

Tetrahedral

A

4 bonding pairs
109.5 degrees

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27
Q

Trigonal bipyramidal

A

5 bonding pairs
90- 120 degrees

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28
Q

Octahedral

A

6 bonding pairs
90 degrees

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29
Q

What are giant covalent structures?

A

they are covalently bonded into a giant lattice structure. Each atom has multiple covalent bonds which give the substance a high boiling point as require a lot of energy to break.

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30
Q

What is diamond?

A

Made of carbon atoms bonded to four further carbon atoms. Rigid tetrahedral structure.

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31
Q

What is thr structure of graphite?

A

Each carbon atom is bonded to three others in flat hexagonal sheets. This means there is 1 delocalised electron per carbon atom which can move freely and allow graphite to conduct electricity. Graphite can therefore be used as an electrode. Intermolecular forces between layers of graphite are weak and can easily slide over eachother meaning it can be a lubricant.

32
Q

What is metallic bonding?

A

Consists of a giant lattice of positively charged ions surrounded by a sea of delocalised electrons. There are very strong electrostatic forces of attraction between oppositely charged particles ( lot of energy to break bonds).
The greater the charge on the pos ion, the stronger the attractive force.
Ions larger in size produce a weaker attraction due to their greater atomic radius.

33
Q

What is metallic structure? Eg Aluminium

A

Often good conductors as delocalised electrons can move and carry a flow of charge.
Metals are also malleable because the uniform layers of pos ions are able to slide over one another.
The electrostatic forces of attraction between pos ions and delocalised electrons are very strong and require a lot of energy to overcome. Therefore metallic substances have high melting points and are nearly always solid at room temp.

34
Q

How are the elements arranged in the periodic table?

A

Increasing proton number.

35
Q

What is the variation of the melting points of the period 2 elements?

A

peaks towards the middle

-lithium and beryllium have metallic bonding so their melting points increase due to a greater positive charge of the ions. ( Li= +1 and Be= +2). There are more electrons released as free electrons in the beryllium lattice so the attractive electrostatic forces are greater then for lithium.

-Boron and carbon form giant covalent lattices with very strong covalent bonds which require a lot of energy to break.

-Nitrogen, oxygen, fluorine and neon are all small simple covalent molecules with weak intermolecular forces. Have low melting points.

36
Q

What is the variation of the melting points of the period 3 elements?

A

-Sodium, mg and al are metals with metallic binding. Their melting points increase due to greater positive charged ions (Na =+1, Mg= +2, Al= +3). More electrons are released as free electrons so the attractive electrostatic forces release from Na to Al).
- Silicon have a strong covalent structure
- phosphorous, sulfur and chlorine have weak intermolecular forces.
-argon is a gas at room temp with a very low melting point as it’s a very stable atom with a full outer shell of electrons

37
Q

Nitrate ion

A

NO3-

38
Q

Sulfate ion

A

(SO4)2-

39
Q

Carbonate ion

A

(CO3)2-

40
Q

Ammonium ion

A

NH4+

41
Q

Bircarbonate ion

A

(HCO3)-

42
Q

Copper ion

A

Cu2+

43
Q

Zinc ion

A

Zn2+

44
Q

Zinc ion

A

Zn2+

45
Q

Lead ion

A

Pb2+

46
Q

Iron (II) ion

A

Fe2+

47
Q

Iron (III) ion

A

Fe3+

48
Q

Group 2 reactions with water

A

Group 2 metals react with water in a redox reaction to produce a metal hydroxide and hydrogen. The metal hydroxide forms an alkaline solution.
Magnesium reacts very slowly with liquid water but the reaction is much faster with steam as it provides extra energy.
Down the group, the reactivity of the group 2 metals increases. Calcium, strontium and barium all react with cold water with increasingly vigorous reactions.

49
Q

What happens when steam is used to react water and magnesium?

A

The magnesium burns with a bright white flame to form hydrogen and magnesium oxide ( a white powder).

50
Q

How does the metal oxides of the group 2 elements react with water?

A

To form basic metal hydroxides. The alkalinity of these increases down the group as the solubility of the metal hydroxides increases.

51
Q

Group 2 compounds reaction with oxygen

A

Group 2 metals react with oxygen to form oxides. Once the reaction has been initiated it is vigorous. Strontium and barium can react with excess oxygen and heat energy to form metal peroxides.

52
Q

Thermal decomposition of group 2 carbonates

A

Group 2 carbonates become more stable to heat as you go down the group. This means as you go down the group, the carbonates have to be heated more strongly for thermal decomposition to occur. The smaller 2+ ion has a higher charge density than a larger 2+ ion. An increasing charge density means the ion will have a distorting effect on neighbouring negative ions. The positive ion attracts the delocalised electrons on an oxygen atom in the carbonate ion towards itself. The carbonate ion becomes polarised which weakens the C-O bond meaning with some heat this easily breaks to release CO2.

53
Q

What’s the general equation for the thermal decomposition of the group 2 carbonates?

A

XCO3 -> XO + CO2

Carbonate and oxide are white solids and carbon dioxide is a colourless gas.

54
Q

Solubility of group 2 carbonates

A

Decreases down the group

55
Q

Solubility of group 2 hydroxides

A

Increases down the group. Magnesium hydroxide is the least soluble so is used in medicine as an antacid as is alkaline and can neutralise acids. Calcium hydroxide is used in agriculture to neutralise acidic soils.

56
Q

What is first ionisation energy?

A

The minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state. Measured in kjmol-1.

57
Q

First ionisation energy of Na

A

Na(g)-> Na+ (g) + e-

58
Q

How do successive ionisation enthalpies occur?

A

When further electrons are removed. Requires more energy because as electrons are removed the electrostatic force of attraction between the positive nucleus and the negative outer electron increases. Ionisation energy increases to overcome this attraction.

59
Q

First ionisation energy along a period

A

Increases due to a decreasing atomic radius and greater electrostatic forces of attraction

60
Q

First ionisation energy down a group

A

Decreasing due to an increasing atomic radius and electron shielding

61
Q

What do successive ionisation energies involve?

A

Removing 1 mole of electrons from one 1 mole of gaseous ions.

62
Q

Why do successive ionisation energies increase?

A

Atomic radius decreases and there is greater attraction between outer shell electrons and the nucleus.

63
Q

What is the equation for change of energy?

A

Plancks constant x frequency of light

64
Q

What is the formula for change of energy?

A

Plancks constant x frequency of light

65
Q

What is the equation for speed of light?

A

Wavelength x frequency

66
Q

What is the visible light region of the EM spectrum in order of increasing frequency and decreasing wavelength?

A

Infrared, visible, UV

67
Q

Emission

A

Atoms can absorb electromagnetic radiation of the correct energy and promote electrons to a higher energy level. As the electrons fall back down to their ground state they release electromagnetic radiation. The light emitted has different wavelengths (emission spectrum).

68
Q

Absorption

A

If a photon with an energy the same as the gap between two energy levels hits an atom, then the electron in the lower energy level can absorb the energy and be promoted to the higher energy level. The photon will not be absorbed if its energy does not correspond to the difference between two energy levels.
To

69
Q

How to obtain an absorption spectra?

A

White light is shined on the material that is being investigated. White light is used as it contains all the wavelengths of visible light.

70
Q

What are similarities between the emission and absorption spectra?

A

Lines become closer at higher frequencies. Lines are in the same position for a given element.

71
Q

What are differences between absorption and emission spectrums?

A

Emission spectrums consist of bright coloured lines on a black background.
Absorption spectrum consist of black lines on a bright background.

72
Q

Lithium ion flame test

A

Red
(When you lie u turn red)

73
Q

Sodium ion flame test

A

Yellow

74
Q

Potassium ion flame test

A

Lilac

75
Q

Calcium ion flame test

A

Orange-red

76
Q

Barium ion flame test

A

Green

77
Q

Copper ion flame test

A

Blue green

chemistry (7 decks)

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