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A levels > Chemistry AS - ABG > Flashcards

Chemistry AS - ABG Flashcards

0
Q

What is an element?

A

A substance that cannot be broken down chemically into anything simpler

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1
Q

What is an atom?

A

The smallest possible particle of an element

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2
Q

What is a compound?

A

A substance composed of 2 or more elements chemically bonded together

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3
Q

What is an ion?

A

A particle formed by an atom or group of atoms losing or gaining electrons

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4
Q

What is a molecule?

A

An arrangement of atoms covalently bonded together

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5
Q

What is a covalent bond?

A

A shared pair of electrons holding atoms together in a molecule

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6
Q

What is an ionic bond?

A

Electrostatic attraction between oppositely charged ions

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7
Q

What is the relative atomic mass of an element?

A

The average mass of an atom of an element on a scale where an atom of carbon-12 12

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8
Q

What is the relative isotopic mass of a substance?

A

The mass or an atom of an isotope of an element of a scale where an atom of carbon-12 is 12.

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9
Q

What is the relative molecular/formula mass of a substance?

A

The average mass of a molecule or formula unit on a scale where an atom of carbon-12 is 12

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10
Q

What is a Mole

A

An amount of a substance that contains exactly the same number of particles as there are atoms in 12g of C-12

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11
Q

What is stoichiometry?

A

The relative number of moles of each substance that react together

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12
Q

What is the molecular formula of a substance?

A

Shows the number of of atoms of each element in one molecule of a substance

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13
Q

What is the empirical formula of a substance?

A

Shows the simplest whole number ratio of atoms of elements present in a substance

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14
Q

What is a salt?

A

A compound formed when the H+ ion from an acid is replaced by a metal or ammonium ion

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15
Q

What are acids?

A

Proton donors - produce H+ ions in aqueous solution

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16
Q

What are bases?

A

Proton acceptors - remove H+ ions In Aqueous solution

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17
Q

Name the 2 main indicators used in titrations, and give observations

A

1) Methyl orange - turns yellow to red when acid is added to an alkali 2) Phenolphthalein - red to colourless when acid is added to alkali

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18
Q

What is spin pairing?

A

Electrons in orbitals spinning in different direction

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19
Q

What shape is a s-orbital?

A

Spherical

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20
Q

What shape is a p-orbital?

A

Dumbbell

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21
Q

How many p orbitals are there in one sub-shell?

A

3, 2 electrons in each, orbitals at right angles to each other

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22
Q

What is first ionisation energy?

A

The energy needed to remove one electron from one mole of gaseous atoms to form one mole of gaseous 1+ ions

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23
Q

How does nuclear charge affect ionisation energy?

A

More protons in nucleus -> more positively charged nucleus -> stronger attraction for electrons -> higher IE

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24
Q

How does distance from nucleus affect ionisation energy?

A

Attraction falls off rapidly with distance -> e- close to nucleus -> much more strongly attracted -> higher IE

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25
Q

How does shielding affect ionisation energy

A

As the number of electrons between outer electrons and nucleus increases, shielding increases, so attraction between outer and electrons decreases -> IE is lower

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26
Q

What is second ionisation energy?

A

Energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions

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27
Q

Why do successive ionisation energies increase within each shell?

A

Electrons are being removed from an INCREASINGLY POSITIVE ION - LESS REPULSION amongst remaining electrons, so the are HELD MORE STRONGLY by the nucleus

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28
Q

When do big jumps in successive ionisation energies increase, and why?

A

When a new shell is broken into - electron is being removed from a shell closer to the nucleus, and therefore is more strongly attracted to the nucleus

29
Q

Why do ionic compounds only conduct electricity when molten or dissolved?

A

Ions in a liquid are free to move and carry a charge When solid they are held in place by strong ionic bonds

30
Q

Why do ionic compounds have high melting points?

A

The giant ionic lattices are held together by strong electrostatic forces Lots of energy is needed to overcome these forces and break the ionic bonds

31
Q

Why are ionic compounds soluble in water?

A

Water is polar, and attracts the ions away from the ion, forming ion-dipole complexes

32
Q

What shape do the atoms in a diamond crystal lattice (covalent) have?

A

Tetrahedral

33
Q

Which properties of diamond can be attributed to its strong covalent bonds?

A

Very high melting point - sublimes Hard Good thermal conductor - vibrations travel easily Won’t conduct electricity - all outer electrons are held in localised bonds Won’t dissolve in any solvent

34
Q

What are allotropes?

A

Different forms of the same element in the same state e.g. Diamond and Graphite

35
Q

Describe the structure of graphite

A

Each carbon has 3 covalent bonds, forming a sheet of flat hexagons each carbons’s 4th electron is delocalised between the hexagon sheets the sheets are bonded together by weak Van Der Waal’s forces

36
Q

How does graphites structure give rise to its different chemical properties?

A

1) weak bonds between layers of graphite are easily broken - slippery, used as a dry lubricant, used in pencils 2) delocalised electrons are free to move & conduct an electric current 3) Far apart layers - less dense than diamond - makes strong, lightweight sports equipment 4) Strong covalent bonds - very high melting point - insoluble

37
Q

describe giant metallic lattice structures:

A

1) delocalised electrons in metal atoms outer shell - electrons free to move - positive metal ion left 2) +ve metal ions attracted to -ve metal ions - form lattice of +ve ions in sea of -ve electrons - metallic bonding

38
Q

How does metallic bonding explain metals properties?

A

1) Number of delocalised electrons per atom affects mpt: - more d.e, stronger bonding, higher mpt - (size of metal ion and lattice structure also affect mpt) 2) no bonds holding ions together, metal ions can slide past eachother - metals are malleable and ductile 3) delocalised electrons pass kinetic energy to each other - good thermal conductors 4) delocalised electrons carry a current - good electrical conductors 5) Strong metallic bonds - insoluble except liquid metals

39
Q

How does electronegativity make a bond polar?

A

bonding electrons are pulled towards the more electronegative atom the greater the difference in electronegativity, the more polar the bond

40
Q

what are the 3 types of intermolecular forces?

A

1) Hydrogen bonding (strongest) 2) Permanent dipole-dipole interactions 3) Van der Waals forces

41
Q

Why does reactivity increase down group 2?

A

As you go down the group: 1) Ionisation Energies decrease - Increasing atomic radius and shielding 2) when grp 2 elements react they lose electrons - lower I.E. = easier to lose electrons = more reactive

42
Q

What is produced when group 2 metals react with water? give an equation

A

Metal hydroxide formed M (s) + 2(H2O) (l) -> M(OH)2 (aq) + H2 (g)

43
Q

What is produced when group 2 metal Oxides react in water?

A

Aqueous metal hydroxide ions MO (s) + H2O (l) -> (M)2+ + 2OH- (aq)

44
Q

why do group 2 Oxides form more alkaline solutions as you go down the group?

A

the hydroxides become more soluble

45
Q

What is formed when group 2 elements burn in oxygen?

A

solid white oxides form 2M + O2 -> 2MO

46
Q

what is thermal decomposition?

A

when a substance breaks down when heated

47
Q

give an equation for the thermal decomposition of a group 2 metal

A

MCO3 -> MO (s) + CO2

48
Q

how does thermal stability change down group 2?

A

thermal stability increases - more heat is required to break it down

49
Q

give examples of uses for group 2 metals

A

1) Ca(OH)2 - slaked lime - used in agriculture to neutralise acid soils 2) Mg(OH)2 -antiacid - indigestion

50
Q

give the equation for neutralisation

A

H+ + OH- -> H2O

51
Q

what colour and state is fluorine at rtp?

A

yellow gas

52
Q

what colour and state is chlorine at rtp?

A

green gas

53
Q

what colour and state is bromine at rtp?

A

red-brown liquid

54
Q

what colour and state is Iodine at rtp?

A

grey solid

55
Q

Describe and explain the pattern of change in mpt/bpt as you go down group 7?

A

1) Boiling/ melting points increase down group 7 2) size and relative mass of atoms increase: = strength of VDW’s increases

56
Q

describe and explain the change in reactivity as you go down group 7

A

Less reactive as you go down group 7 1) halogens react by gaining an electron - reduced -> oxidising agents 2) as you go down group 7, atom size increases -> outer electrons further from nucleus 3) outer electrons are shielded from increasingly positive nucleus - harder for larger halogen atoms to attract an electron 4) larger atoms = less reactive = less oxidising

57
Q

What is the general ionic equation for reacting Silver Nitrate with a Halogen?

A

Ag+ (aq) + X- (aq) -> AgX (s)

58
Q

describe the method for testing for halides:

A

1) Add Nitric acid to halide solution - removes unwanted ions 2) add AgNO3 (aq) 3) a coloured precipitate forms 4) add NH3

59
Q

describe the observations when the following halide ions undergo the test: 1) Cl- 2) Br- 3) I-

A

1) Cl- : White ppt, dissolves in dilute NH3 2) Br- : Cream ppt, dissolves in conc. NH3 3) I - : yellow ppt, insoluble in NH3

60
Q

What is Sodium Chlorate (I) (NaClO) used for?

A

Bleach - water treatment, bleaching paper and textiles

61
Q

What is the equation for the formation of bleach? what conditions are needed?

A

2NaOH (aq) + Cl2 (aq) -> NaClO (aq) + H2) (l) rtp

62
Q

What is the equation for the disproportionation of chlorine in water

A

Cl2 (g) + H2O (l) <–> HCl (aq) + HCLO ( Chloric (i) acid)

63
Q

Give the equation of the ionisation of Chloric (I) acid:

A

HClO (aq) + H2O (l) <–> ClO- (aq) + H3O+

64
Q

Why is chlorine used in water treatment?

A

1) Kills disease-causing microorganisms 2) prevents reinfection in water supply 3) prevents growth of algae - eliminating bad smells and tastes - removes discolouration

65
Q

What are the disadvantages of adding chlorine

A

1) Chlorine gas is very harmful - irritates respiratory system 2) Liquid chlorine causes severe chemical burns 3) Chlorine reacting with organic compounds makes Chlorinated Hydrocarbons - carcinogens

66
Q

Why is fluoride ions added to water in some parts of the uk, and why are some people concerned by it?

A

1) Fluoride ions prevent tooth decay 2) small amount of evidence linking fluoride to bone cancer 3) toothpaste is already fluorinated

67
Q

On what occasions can Hydrogen Bonding occur, and Why?

A

1) When H is covalently bonded to F, N or O 2) H has a high charge density (v. small atom) , others are very electronegative 3) Bond is polarised 4) H of one molecule can form a weak bond with F,N or O on another molecule

68
Q

How does hydrogen bonding affect the physical properties of substances, compared with non polar molecules of similar sizes?

A

Hydrogen bonding makes molecules: 1) Soluble in water 2) Higher Boiling and Melting points

69
Q

How are molecules of H2O held together in ice?

A

by H bonds, in a lattice

70
Q

Why is ice less dense than water?

A

H bonds are relatively long

A levels (12 decks)

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