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Summer exams > Chemistry 2 > Flashcards

Chemistry 2 Flashcards

1
Q

Elements are composed of atoms which may be linked to form

A

molecules

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2
Q

relative atomic mass Ar

A

the mass of 1 atom relative to 1/12 the mass of 1 atom of carbon-12

carbon-12 is the most common isotope of carbon and its mass is 12

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3
Q

Relative MOLECULAR mass Mr definition

A

-the mass of 1 molecule of an element or compound compared with 1/12th the mass of 1 atom of carbon-12

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4
Q

Relative MOLECULAR mass Mr HOW TO CALCULATE

A

addition of Ar atoms in molecule

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5
Q

the mole

A

A mole of carbon 12 Is the amount of chemical which contains exactly 12g of 12C

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6
Q

The Avogadro constant and what it is

A

12 /1.9926 x 10^-23 =6.022 x 10^23

1.9 is the mass of one atom

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7
Q

number of moles =

A

mass of 1 substance/mass of 1 mole (g)

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8
Q

molarity of solutions

A

the amount of solute in moles in 1 dm^3 of solution

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9
Q

calculating molarity of solutions

A

-convert to dm^3 for example if 10cm of 0.1 then 10/1000 x 0.1 is how many in dm^3

-then multiply by Mr

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10
Q

A large Kc value indicates

A

the product concentration is high and the forward reaction is favoured.

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11
Q

le chetiliers

A

poster

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12
Q

haber process

A

iron can be used as catalyst

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13
Q

weak acids and bases only

A

partially dissolve in water

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14
Q

The pH of a weak acid depends on two factors

A

-the concentration of the acid
-its pKa value (pKa = -log10 Ka)

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15
Q

to calculate the pH of weak acids …

A

-use the Ka formula to find concentration of [H]
-then put in log formula

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16
Q

strong bases

A

-fully ionised
-concentration of hydroxide (OH-) is equal to the concentration of the base

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17
Q

weak bases

A

-only partially ionised
-equlibrium is established

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18
Q

aqueous ammonia is the most common

A

weak base

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19
Q

the base ionisation constant

A

Kb = [BH+]
look at slide 14 on Acids and bases part 2

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20
Q

Henderson-hasselbalch equation

A

-possible to calculate the pH of a buffer solution

pH = pKa + log10 ([base]/[acid])

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21
Q

acid =

A

a substance which contains hydrogen and releases hydrogen ions (H+) when dissolved in water

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22
Q

base =

A

substance which reacts with an acid to form a salt plus water

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23
Q

acid is proton

A

donor

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24
Q

base is proton

A

acceptor

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25
Q

acid + base =

A

salt + water

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26
Q

how the atom arose

A

look on ppt atomic structure

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27
Q

atomic number

A

equal to the number of positive charges or protons in the atom

bottom left

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28
Q

labelled numbers

A

-atomic mass top = protons + neutrons
atomic number = protons + electrons unless isotope

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29
Q

if the atomic mass is different

A

number of neutrons changes

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30
Q

the 1st ionisation energy

A

defined as the energy required to turn 1 mole of an element in the gas state into 1 mole of its ions

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31
Q

patterns in ionisation energy increasing in reactivity across the periodic table

A

as you go up the table increase in charge and attraction

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32
Q

isotopes are elements containing atoms with the same

A

atomic number

but different atomic mass

they can be radioactive

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33
Q

intra molecular bonds

A

ionic

covalent

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34
Q

inter-molecular bonds

A

dispersion

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35
Q

a vertical column called a group contains elements with

A

similar chemical properties due to same number of valence electrons

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36
Q

horizontal row is called a

A

period
atomic number increases from left to right along the row

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37
Q

groups

A

-have same number of outer shell electrons

-down a group = increasing number of electron shells
-melting and boiling points decreases in groups 1 to 4 and increase In groups 5 to 8

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38
Q

atoms become larger

A

-group 1 and 2 elements become more reactive
-groups 6 and 7 elements become less

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39
Q

within periods they have the same number of electron

A

shells

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40
Q

across a period

A

-elements become more electronegative
-atomic radius decreases
-melting and boiling points rise to a peak (group 4) then fall

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41
Q

electronegativity increases across the periodic table

A

across and up

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42
Q

after passing an electron from the metal to non both adopt a

A

noble gas configuration

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43
Q

ionic bonding requires electron

A

transfer

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44
Q

another example of ionic bonds

A

ionic bonds can be pH sensitive and not all ionic bonds can be removed by electrolysis

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45
Q

covalent bonding

A

-incomplete electron transfer leads to a shared electron noble gas configuration

each atom donates an equal number of electrons

46
Q

coordinate bonding

A

sharing electrons only supplied by one side

47
Q

metallic bonding

A

-metal elements can share their outer electrons

48
Q

to infer a structure

A

notice the electrons in the outer electron shell

-repulsion axis

49
Q

intermolecular forces

A

-hydrogen bonds, electrostatic bonds
-van der Waals bonds and teh hydrophobic exclusion

50
Q

intermolecular forces strings of interactions in order

A

-covalent bond
-ionic
-hydrogen
-van der waals

51
Q

dotted line

A

hydrogen bond

52
Q

filled in line

A

covalent

53
Q

electrostatic dispersion bonds

A

-dipole-dipole
-dipole-ion

beyond the hydrogen

54
Q

molecular orbitals

A

shapes of molecules explained by atomic and molecular orbital theory

55
Q

Stereochemical formulae different lines mean

-solid line

A

bond in the plane of the paper

56
Q

Stereochemical formulae different lines mean

-dashed line/diminishing wedge or a striped wedge

A

bond that goes behind the paper

57
Q

Stereochemical formulae different lines mean

-enlarging wedge

A

bond that goes in front of the paper

58
Q

a single line represents

A

single covalent bond

59
Q

the ends of all bonds are occupied by

A

a hydrogen atom

60
Q

saturated molecules

A

contain single c - c bonds

61
Q

Alkane formula

A

Cn H(2n +2)

62
Q

Fractional distillation

A

different fractions obtained by fractional distillation crude oil

63
Q

catalytic cracking

A

of longer chain alkanes is carried out as a means of increasing the yield of the petrol fraction obtained in fractional distillation

64
Q

double bonds

A

made up of two electron pairs

65
Q

Alkenes

A

contain one c to c double bond these carbon atoms are sp^2 hybridised

66
Q

ethene no rotation around

A

the c= c bond

67
Q

alkene different structures on ppt

A

organic 1 FA slide 26

68
Q

Alkynes

A

ethyne has a title bond and formula C2H2

69
Q

aldehydes and ketones

A

both carbonyl compounds

70
Q

Aldehydes are easily oxidised to

A

carboxylic acids

71
Q

esters

A

produced by mixing together an acid and an alcohol

72
Q

triglycerides and fat

A

made by esterification of glycerol a special type of alcohol with three fatty acids

glycerol forms polar head of these special esters

three fatty acids with long alkane chains join this to form saturated fats

73
Q

trans-fats have one or more

A

double bonds in their fatty acids which allow each of the side chains to line up parallel so they pack well

74
Q

cis-fats

A

have tails sticking out at angle so poorer packing results in lower melting point

75
Q

naming the molecules ppt

A

in organic 1 on page 37

76
Q

Alkanes

A

unreactive non-polar hydrocarbons

-strong c-c bonds
-no dipole to attract charged molecules

reactions restricted
-combustion
-substituition reactions in which they rely on highly reactive free radicals

77
Q

Alkenes

A

-much more reactive than alkanes
-due to presence of double bond
-unsaturated alkenes turn to saturated molecules following addition reactions

78
Q

electron rich pi bond attracts

A

molecules with sigma+ charge

79
Q

benzene and electrophilic substituition

A

-instead of electrophilic addition reactions benzene undergoes electrophilic substituition
-positive ion or group of atoms is attracted to the pi electron cloud

-H atom removed

80
Q

Alcohol general formula

A

CnH2n+1 - OH

81
Q

primary alcohol

A

has 2 hydrogen atoms attached to the carbon atom OH group

82
Q

secondary alcohol

A

has 1 hydrogen atom attached to the carbon atom that bears the OH group

83
Q

tertiary alcohols

A

no single hydrogen atoms on the carbon atom that bears the OH group

84
Q

oxidation of alcohols

A

strong oxidising agent required in a reaction that can be used in breathalyser tests for ethanol

85
Q

oxidation of ethanol to ___ then to

A

ethanal then to ethanoic acid

86
Q

carbonyl compounds the two main types

A

aldehydes and ketones

87
Q

aldehydes vs ketones

A

oxidised by weak oxidising agents such as Benedict’s reagent to carboxylic acids
-blue copper two ions are reduced to a brick red copper oxide precipitate

-ketones do not reduce benedicts
-glucose has aldehyde groups so acts as a reducing sugar on benedicts

88
Q

the carbonyl group is attractive to

A

nucleophiles

89
Q

carboxylic acids

A

-soluble in water
-form h bonds between two acid molecules or between acid molecules and water
-do not dissociate completely therefore ARE WEAK ACIDS

90
Q

fats are

A

triesters

91
Q

prostaglandins

A

-affect many physiological processes
-act as hormones or hormone modulatirs
-derived from C20 fatty acids with at least three double bonds

92
Q

spectrometry look on

A

ppts

93
Q

IR spectroscopy

A

-radiation across range is passed through substance
-strong bonds = low mass atoms vibrate and absorb energy at high frequencies
-weak bonds = high atomic masses absorb at low frequencies

94
Q

what can IR spectroscopy tell us

A

-used to identify and differentiate between species of bacteria
-show conversion of prion protein into misfiled pathological form

95
Q

Atomic emission spectroscopy

A

-atomic emission spectroscopy (ICP_AES)
-aerosolised sample is carried into argon gas plasma
-radiation emitted as atoms return to lower energy levels

96
Q

nuclear magnetic reasoning

A

-atomic nuclei with odd mass numbers such as hydrogen, posses a spin which can align with or against an applied magnetic field

97
Q

UV/visible spectroscopy

A

-used to detect electronic transistions between molecular energy levels
-group of atoms producing a characteristic absorption called a chromophore

98
Q

beer-lambert law

A

A = E x C x L

aborbance
mlar absorptivity Lmol-1cm-1 (extinction coefficient)
concentration L-1
path length cm

99
Q

UV/visible spectroscopy

A

-spectra result from electronic transitions between molecular energy levels
-molecular orbitals represent different quantised energy levels for molecules in the same way as atomic orbitals do for atoms

100
Q

type of electron transitions
-sigma

A

represent valence bonds which have the lowest energy levels

101
Q

type of electron transitions
-pi

A

pi bonds are at a higher energy level

102
Q

type of electron transitions
-non bonding electrons

A

N, O Halogen or S

which do not participate in bonding occupy the highest energy ground state

103
Q

benzene rings

A

have delocalised electrons as do heterocyclic rings with carbon and nitrogen

104
Q

beta carotene

A

conjugated because of its alternating double and single bonds this decreases the energy needed for a pi-pi* transition
shifting to a longer visible wavelength

105
Q

what do we use UV for?

A

-locating samples in purification process
-checking concentration of product
-changes in conformation of proteins/protein
-enzyme actvity

106
Q

types of isomerism

A

-structural or stereoisomerism

107
Q

structural isomerism

A

simply molecules with different structural formulae
-branched chains
-position of functional groups
-different functional groups

108
Q

steroisomerism

A

-same structural formulae the orientation of bonds is different
-geometric isomerism
-optical isomerism

109
Q

chain isomers

A

chan molecules can have varying amounts of branches

110
Q

positional isomers

A

functional groups can adopt several different positions

111
Q

E/Z isomerism

A

special case of geometric isomerism when you have 4 different groups around the double bond

-if highest rank groups are together = Z
-opposite = E

112
Q

D and L amino acids

A
  • POSITIVE enantiomers can also be described as d enantiomers (right handed clockwise)
  • NEGATIVE enantiomers can also be described as L enantiomers (left handed anti-clockwise)

Summer exams (6 decks)

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