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Chemistry > Jan 9 > Flashcards

Jan 9 Flashcards

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1
Q

Separating oxygen and nitrogen from air + procedures

A

By fractional distillation
1. Air is purified to remove dust, CO2, water vapour etc.
2. Purified air is cooled to -200˚C by repeating compression and rapid expansion to change to liquid
3. Liquid air at -200˚C is introduced at the lower part of the fractionating column
4. Liquid air is warmed slowly and gases are collected one by one. At -190˚C, nitrogen (bp -196˚C) vaporises and is separated from the liquid mixture first. Argon (bp -186˚C) vaporises next and oxygen (bp -183˚C) follows at higher temperatures and is then collected.

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2
Q

Test for oxygen

A

relight a glowing splint

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3
Q

Test for hydrogen

A

a “pop” sound is produced by a burning splint

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4
Q

Saturated solution

A

Saturated solution is when a solution contains the maximum amount of solute at a certain temperature

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5
Q

Filtration

A

separate an insoluble solid from solution

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6
Q

Evaporation

A

separate a dissolved solid with high boiling point from its solution

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7
Q

Crystallisation + procedure

A

separate a dissolved solid from a solution
1. Saturated solution is obtained by heating solution to evaporate some of the solvent
2. Large crystals will form upon slow cooling of saturated solution (solubility of solid in solvent decreases with decreasing temperature)
3. Filter the large crystals from the saturated solution

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8
Q

Decantation

A

separate much denser insoluble solid from liquid

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9
Q

Simple distillation

A

separate dissolved substance with high boiling point from a solvent

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10
Q

Flame test (procedures)

A
  1. Clean platinum wire by dipping the wire into concentrated hydrochloric acid and hold it in a non-luminous flame until no characteristic flame is given out
  2. Moisten clean wire by dipping it into the acid and into the solid sample
  3. Hold the wire in the hottest part of a non-luminous flame and observe the colour of the flame
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11
Q

Flame test (results)

A

Calcium ion: Brick-red flame
Potassium ion: Lilac flame
Sodium ion: golden yellow flame
Copper (II) ion: bluish green
Magnesium ion: Bright white flame (Not for MgCl2, MgSO4, Mg(NO3)2)
Aluminium ion: White flame
Zinc ion: White-blue flame

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12
Q

Test for chloride

A
  1. Dissolve sample in water
  2. Add excess dilute nitric acid
  3. Add aqueous solution of silver nitrate
  4. White precipitate (silver chloride) will be produced
    silver nitrate + calcium chloride —> silver chloride + calcium nitrate
    silver ion + chloride ion —> silver chloride
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13
Q

Test for water

A

turn blue anhydrous cobalt (II) chloride solid to pink (dry cobalt(II) chloride paper) / white anhydrous copper(II) sulphate to blue

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14
Q

Hydrogen usage

A
  1. Hydrochloric acid
  2. NH3
  3. Fuel (highly flammable —> provides a lot of heat energy, product after burning does not cause pollution)
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15
Q

Chlorine usage

A
  1. Hydrochloric acid
  2. Bleaching solution
  3. Sterilising drinking water / pool
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16
Q

Sodium hydroxide usage

A
  1. Soap
  2. Paper
  3. Bleaching solution
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17
Q

Mineral

A

Mineral is a naturally occurring inorganic solid which has a characteristic chemical composition and highly ordered crystalline structure.

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18
Q

Limestone usage (CaCO3)

A
  1. Construction of buildings and roads
  2. Making glass by heating with sand and sodium carbonate
  3. Neutralising sulphur dioxide in flue gas from power station
  4. Making cement
  5. Neutralising acidity in soil and lakes
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19
Q

Quicklime usage (CaO)

A
  1. Making steel from iron
  2. Neutralising acidity in soil
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20
Q

Slaked lime usage (Ca(OH)2)

A
  1. Making bleaching powder
  2. Neutralising acidity in soil and lakes
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21
Q

Equations of calcium carbonate

A

Calcium carbonate –> (heat) calcium oxide + carbon dioxide [decomposed]
Calcium oxide + water –> calcium hydroxide (s)
Calcium hydroxide (s) + water –> calcium hydroxide (aq)
Calcium hydroxide + carbon dioxide –> calcium carbonate + water
Calcium carbonate + carbon dioxide + water –> calcium hydrogen carbonate
Calcium carbonate + dilute hydrochloric acid –> calcium chloride + carbon dioxide + water [effervescence occur]
Calcium carbonate + nitric acid –> calcium nitrate + carbon dioxide + water

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22
Q

Process of chalk to marble

A

sea animals with shell died
shells fell to sea bed and built up deposit
under high pressure, deposit is converted to chalk after millions of years
under high pressure and temperature, chalk is converted to limestone
under more higher pressure and temperature, limestone is converted to marble

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23
Q

Acid rain equation

A

water + carbon dioxide –> carbonic acid
water + nitrogen oxide –> nitric acid
water + sulphur dioxide –> sulphurous acid

calcium carbonate + carbonic acid –> calcium hydrogen carbonate solution

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24
Q

Atom

A

Atom is the smallest particle of an element that still has its chemical properties and cannot be split into simpler particles by chemical methods.

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25
Q

Element

A

Element is a pure substances which cannot be broken down into anything simpler by chemical methods.

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26
Q

Compound

A

Compound is a pure substance that contains two or more different kinds of element that is chemically bonded. Its formation is a chemical change in which new substances are formed. It can be separated into its constituent by chemical methods.

27
Q

Mixture

A

Mixture is a substance that contains two or more different kinds of substances that are not chemically bonded. Its formation is a physical change in which no new substances are formed. It can be separated into its constituents by physical methods.

28
Q

Isotope

A

Isotopes are different atoms of the same element that have the same number of protons but a different number of neutrons

29
Q

Relative abundance

A

Relative abundance is the distribution of isotopes in nature of an element.

30
Q

Relative isotopic mass

A

Relative isotopic mass of a particular isotope of an element is the mass of an atom of the isotope relative to the 1/12 of the mass of a carbon-12 atom.

31
Q

Relative atomic mass

A

Relative atomic mass of an element is the average isotopic mass of all the natural isotopes on the 12C=12.00 scale.

32
Q

Same period properties

A
  1. Atomic size increases from left to right as the attraction between electron and the nucleus increases
  2. Group I to III mp and bp increases as metallic bond strength increases with increasing number of outermost electrons
    Group IV mp and bp is maximum as its element has giant covalent structure
    Group V to VII mp and bp is low as they have weak van Der Waals’ forces exist between molecules
    Group 0 mp and bp are low because they are all monatomic molecules, weak van Der Waals’ forces exist between molecules
  3. ability to attract electrons increases across the period as the number of protons increases across the period which causes increased attraction between incoming electron and nucleus.
33
Q

Same group properties

A
  1. Atomic size increases down the group as the number of occupied electron cells increases down the group
  2. Group I to II mp and bp decreases down the group as atomic size increases down the group, attraction between cations and delocalised electrons in the metal decreases which leads to weaker metallic bond.
    Group VII to 0 increases down the group as molecular size increases down the group, leads to increasing length of van Der Waals’ forces.
  3. Ability to attract electrons decreases down the group as the number of occupied electron shells increases down the group which there is a weaker attraction between the incoming electrons and the nucleus.
34
Q

Group I (physical + chemical properties)

A

alkali metals
reactivity increases down the group
soft and can be cut by a knife
shiny when freshly cut but tarnish easily
low density
low mp and bp
react with oxygen to form metal oxide
react with water to form metal hydroxide (Li fizzes steadily on water surface, Na melts to form a silvery ball and fizzes quickly on water surface, K melts to form silvery ball and fizzes furiously on water surface, hydrogen gas catches fire and produce a lilac flame
react with chlorine to form metal salts

35
Q

Group II (physical + chemical properties)

A

alkaline earth metals
reactivity increases down the group (less reactive than alkali metals as they need to lose more electrons than alkali metals)
soft but harder + denser than alkali metals
shiny when freshly cut and tarnish easily
high mp or bp
reacts with oxygen to form metal oxide when heated in air
reacts with dilute hydrochloric acid to form metal chloride and hydrogen gas (metal dissolves and gas bubbles are evolved)
reacts with chlorine to form metal salts
some reacts with water to form metal hydroxide + hydrogen (Mg reacts very slowly with water, Ca reacts readily and gives steady stream of gas bubbles)

36
Q

Group VII (physical + chemical properties)

A

halogens
reactivity decreases down the group
mp and bp increases down the group
F is pale yellow gas
Cl is greenish yellow gas
Br is reddish brown liquid (vaporise easily to give orange vapour)
I is black solid (sublimes easily to form purple vapour)
At is black solid
poor conductors of electricity
colours are darker down the group
less volatile down the group
reacts with hydrogen to give hydrogen halides

37
Q

Group 0 (physical + chemical properties)

A

noble gases
stable because it fulfils octet rule / duplet rule
reactivity decreases down the group
colourless
low mp and bp
very low density
poor conductors of heat and electricity
reacts with fluorine under special condition (Kr, Xe)

38
Q

Neon usage

A

advertising signs as it produces a distinctive red-orange light when electricity is passed through it

39
Q

Argon usage

A

added to the space above wine in wine barrels to stop oxygen in air to react with wine as argon is denser than air
fill tungsten light bulbs because it does not react with the hot tungsten filament

40
Q

Monatomic ion

A

Monatomic ion is an atom that loses or gains one or more electrons

41
Q

Polyatomic ions

A

Polyatomic ions are ions that are formed from a group of atoms

42
Q

Conductor

A

Conductors are substances which conduct electricity but are not chemically changed during electrical conduction

43
Q

Electrolyte

A

Electrolytes are substances which conduct electricity in molten state or aqueous solution but not in solid state and are decomposed during conduction.

44
Q

Non-conductor

A

Non-conductors are substances which do not conduct electricity in solid, molten state or aqueous solution

45
Q

Ionic bond

A

Ionic bond is the strong electrostatic forces of attraction between oppositely charged ions. An ionic bond is formed when electrons are transferred from one atom to another. When metal reacts with non-metal to combine to form an ionic compound. Electrons are transferred from atoms of the metal to the atoms of the non-metals.

46
Q

Arrangement name of ionic bond

A

lattice

47
Q

Which 2 compounds have ionic bonds but are not formed by oppositely charged ions

A

ammonium chloride
ammonium sulphate

48
Q

Metallic bond

A

Metallic bond is the strong electrostatic force of attraction between the positive metal ions and the “sea” of delocalised electrons.

49
Q

Covalent bond

A

Covalent bond is the strong electrostatic force of attraction between positively charged nuclei of two bonded atoms and the negatively charged shared electrons. Covalent bond forms when two atoms share a pair of electrons.

50
Q

Diatomic molecules

A

Hydrogen
Oxygen
Chlorine
Nitrogen

51
Q

Lone pair of electrons

A

Lone pair of electrons is a pair of electrons in the outermost shell of one of the atoms in a molecule which is not involved in the bonding.

52
Q

Dative covalent bond

A

Dative covalent bond is a type of covalent bond in which lone pair of electrons come from the same atom.

53
Q

Ionic compounds properties

A
  1. usually in crystalline form
  2. high mp and bp as ions are held by strong ionic bonds
  3. solid compounds do not conduct electricity because the ions are not mobile
  4. molten state or aqueous solution compounds conduct electricity because ions are mobile
  5. hard because ions are held together closely by strong ionic bonds
  6. brittle because ions are attracted by strong ionic bonds. when a force is applied, ions of the same charge may come together and cause repulsion. solid breaks apart
  7. high density because ions are held by strong ionic bond
  8. soluble in water but insoluble in non-aqueous solvents
54
Q

Substance to dissolve in a solvent

A

Attractive force between its particles and solvent particles are strong enough to overcome attractive forces between particles in pure substances. particles in substance must separate from each other and surrounded by solvent particles.

55
Q

Allotrope

A

Allotropes are two or more forms of the same element in which the atoms or molecules are arranged in different ways.

56
Q

Name elements that have giant covalent structures

A

Diamond, Graphite, Silicon dioxide, Silicon

57
Q

Diamond structure

A

a carbon atom is tetrahedrally bonded to four other carbon atoms
covalent bonds are extended in all directions throughout the structure

58
Q

Graphite structure + property

A

layered structure
carbon atoms in the same layer are held together by strong covalent bonds
weak van Der Waals’ forces exist between the layers
can conduct electricity in the graphene plane as there is delocalised electrons in the outermost shell
can be used as making because as layers of graphite are only attracted by weak van Der Waals’ forces so that the layers can slide past each other easily

59
Q

Giant covalent structures properties

A
  1. High mp and bp because atoms are held together by strong covalent bonds
  2. Do not conduct electricity because there are neither mobile ions or delocalised electrons (except graphite)
  3. Very hard because atoms are held together by strong covalent bonds to form a giant network
  4. Insoluble in all solvents because atoms are held together by strong covalent bonds
60
Q

Simple molecular structures properties

A
  1. low mp and bp because molecules are attracted by weak van Der Waals’ forces only
  2. larger molecular size has higher melting and boiling point because the strength of van Der Waals’ forces increases as the molecular size increases.
  3. Do not conduct electricity in both solid and molten
  4. Usually soft because molecules are attracted by weak van Der Waals’ forces
  5. Usually soluble in non-aqueous solvent and slightly soluble or insoluble in water
61
Q

Giant metallic structure properties

A
  1. strength of metallic bond increases from left to right as number of outermost electrons increases.
  2. strength of metallic bond decreases as the atomic size increase down the group, attraction between protons in the nucleus and delocalised electrons decreases down the group
62
Q

Roasting

A

metal sulphide is heated strongly in the presence of excess air

63
Q

Calcination

A

carbonate ores is heated strongly in the absence or limited supply of air

64
Q

Blast furnace

A

Hematite, coke, limestone is added through the top of blast furnace
Hot air is blasted near the bottom of furnace
Coke burns in air to form carbon dioxide
More coke is added (carbon dioxide + carbon –> carbon monoxide)
Carbon monoxide acts as reducing agent and reduces haematite
Molten iron contains impurities trickling downwards in the blast furnace. Impurities are removed using limestone.
Calcium carbonate decomposes under higher temperature (calcium carbonate –> calcium oxide + carbon dioxide)
Calcium oxide reacts with impurities to form calcium silicate (from silicon dioxide). Calcium silicate forms a slag which runs down the furnace and floats on top of molten iron.
Waste gases (nitrogen and oxide of carbon) escape from top of furnace.

Chemistry (3 decks)

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