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Chemistry SL > Chapter 1, 2, & 3 > Flashcards

Chapter 1, 2, & 3 Flashcards

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1
Q

define enrichment

A

separation of isotopes

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2
Q

why is fluorine weird?

A

it only has 1 naturally occurring isotope

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3
Q

what is the Bohr model?

A
  • electrons can occupy DISCRETE energy levels within atom
  • can transition b/w these energy levels by ABSORBING or EMITTING EXACTT amounts of energy
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4
Q

define atomic orbital

A

region of space where there is a high probability of finding an electron

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5
Q

Pauli exclusion principle says . . .

A

2 electrons can only occupy the same atomic orbital if they have OPPOSITE spins

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6
Q

Aufbau principle says . . .

A

electrons fill atomic orbitals of lowest energy first - atomic orbitals within a sublevel (e.g. the 3 orbitals in 2p sublevel) are DEGENERATE

exceptions are Cr & Cu

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7
Q

define degenerate

A

equal energy

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8
Q

which has higher energy, 4s or 3d?

A

3d

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9
Q

where are electrons removed from first?

A

4s orbital is further away from the nucleus than 3d so electrons are removed from the furthest shell first

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10
Q

Hunds rule says . . .

A

electrons fill orbitals in same sub-level singly before pairing up

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11
Q

define wavelength

A

distance b/w 2 crests in oscillating wave

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12
Q

define frequency

A

of waves that pass a point in 1 second

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13
Q

order of decreasing wavelength (increasing energy)

A

radio>microwaves>IR>VIS>UV>X-rays>Gamma

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14
Q

continuous spectrum?

A

all wavelengths + frequencies of visible light are shown

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14
Q

continuous spectrum?

A

all wavelengths + frequencies of visible light are shown

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15
Q

line spectrum?

A

some wavelengths are missing

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16
Q

difference b/w absorption & emission line spectrum

A

ABSORPTION = ALL COLOURS, black streaks
EMISSION SPECTRUM = BLACK, colorful streaks (black EMU wearing COLORFUL sunglasses)

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17
Q

The longer the arrow in a hydrogen emission spectrum . . .

A

the greater the amount of energy that is emitted

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18
Q

how does an emission line spectrum form?

A
  • high voltage passed through gas
  • electrons become excited
  • transition to higher energy levels
  • as electrons fall back down to lower energy levels, transitions accompanied by emission of energy in the FORM OF PHOTONS
  • the energy released by different elements corresponds to dif wavelengths within the visible light range of EM radiation
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19
Q

Briefly explain how a mass spectrometer works

A
  • used to determine relative atomic masses + organic structure
  • sample is vaporised
  • bombarded w/ high energy electrons
  • produces + ions
  • accelerated in electric field
  • cations DEFLECTED in magnetic field
  • depending on mass / charge ratio
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20
Q

higher mass to charge ratio = ?

A

LESS DEFLECTION

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21
Q

which is deflected more, 79Br2+ or 91Br+ ?

A

79Br2+ since it has the LOWER m/z ratio

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22
Q

Solids?

A
  • particles closely packed
  • strong forces b/w particles
  • vibrate about fixed positions
  • fixed shape + volume
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23
Q

Liquids?

A
  • slightly more spread out
  • weaker forces b/w particles
  • can move past each other
  • take shape of container
  • fixed volume
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24
Q

Gases?

A
  • particles very spread out
  • negligible forces
  • no fixed shape / volume
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25
Q

sublimation

A

from SOLID to gas
e.g. iodine or solid CO2

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26
Q

depostion

A

from GAS to solid

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27
Q

difference b/w boiling & evaporation?

A
  • BOILING = occurs at specific temp & takes place throughout the liquid
  • EVAPORATION = only takes place at surface of liquid and can occur at temp. BELOW the boiling point
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28
Q

mixtures

A

contain 2 or more substances mixed together
- NOT chemically bonded
- mixed in ANY proportion
- can be separated by PHYSICAL meanss

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29
Q

heterogenous vs homogenous mixture

A
  • homogenous = only 1 phase, uniform composition throughout
  • heterogenous = more than 1 phase, non-uniform composition
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30
Q

immiscible

A

describes 2 liquids that form a heterogenous mixture

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31
Q

define average atomic mass (Ar)

A

average of the masses of the isotopes in a naturally occurring sample of the element relative to the mass of 1/12 of an atom of C-12

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32
Q

relative molecular mass

A

the sum of the relative atomic masses of the individual atoms making up a molecule

33
Q

relative formula mass

A

(used for ions) the mass of one formula unit relative to the mass of 1/12 of an atom of C-12

34
Q

define empirical formula

A

the simplest whole number ratio of the elements present in a compound

35
Q

percentage yield = ?

A

(actual yield / theoretical yield) x 100

36
Q

What is Avogadro’s law?

A

equal VOLUMES of gases at the same temperature & pressure contain the same NUMBER OF PARTICLE

37
Q

what is constant in avogadro’s law?

A

pressure
temperature

38
Q

When do gases deviate most from ideal behaviour?

A

at HIGH PRESSURE
at LOW TEMPERATURE

hplt = hajar paints leo’s teeth

39
Q

Properties of an ideal gas

A
  • particles are in CONSTANT, RANDOM, STRAIGHT-LINE MOTION
  • collisions b/w particles of an ideal gas are elastic (total kinetic energy conserved)
  • volume occupied by gas is NEGLIGIBLE relative to vol of container
  • NOT INTERMOLECULAR FORCES
  • average kinetic energy of particles is DIRECTLY proportional to ABSOLUTE TEMP. IN KELVIN
40
Q

Why is value pf PV/nRT LESS THAN 1 at moderately high pressures?

A

because of the effects of intermolecular forces

41
Q

Why is value of PV/nRT GREATER THAN 1 at very high pressures?

A

because of the effects of MOLECULAR VOLUME

42
Q

formula for moles when dealing with molar volume

A

n = V / Mv

Mv = 22.7 dm^3 mol-1

43
Q

Boyle’s law

A

VOLUME is INVERSELY PROPORTIONAL to PRESSURE

(downward CURVE)

44
Q

Charles’s Law

A
  • VOLUME is DIRECTLY PROPORTIONAL to TEMPERATURE in kelvin
45
Q

Gay-Lussac’s Law

A

PRESSURE is DIRECTLY PROPORTIONAL to ASBOLUTE TEMPERATURE

46
Q

Use Pa with ??

A

m^3

47
Q

Use kPa with ??

A

dm^3

48
Q

formula for concentration in ppm?

A

mass of solute (g)) / mass of solution (g) x 10^6

49
Q

1 ppm = ?

A

1 mg/ L

50
Q

difference b/w hydrated & anhydrous?

A

hydrated = contain water of crystallisation
anhydrous = have lost their water of crystallisation

51
Q

Experimental determination of empirical formula

A
  • coil of Mg ribbon placed in pre-weighed crucible
  • mass is recorded
  • crucible placed on a clay triangle
  • heated
  • when Mg ribbon starts to burn the lid is lifted slightly to allow more air to enter
  • heating is continued until all Mg has burned
  • after cooling the crucible, its lid & content are reweighed
52
Q

properties of chemical reactions

A
  • new substances are formed
  • bonds in reactants are broken
  • bonds in products are formed –> energy change
  • fixed relationship b/w # of particles of reactants & products resulting in NO OVERALL change in mass
53
Q

what is a standard solution?

A

a solution of KNOWN concentration

54
Q

where is the burette reading taken from?

A

the bottom of the meniscus

55
Q

define atom economy

A

increasing the conversion efficiency of a chemical process
- a measure of the amount of starting materials that become useful products

56
Q

what does a high atom economy mean?

A

fewer natural resources are used & less waste is created

57
Q

atom economy formula

A

(total mass of DESIRED product) / (total mass of ALL products) x 100

58
Q

Outline the use of radioactive isotopes

A
  • isotopes are radioactive as the nuclei of these atoms break down spontaneously
  • when they break down, these radioisotopes EMIT RADIATION which is dangerous to living things
  • 3 forms of radiation: gamma, alpha, beta
  • radioisotopes are used for: power generation, sterilisation of surgical equipment, crime detection, preservation of food

14-C used for carbon dating
isotopes of Iodine used as tracers in medicine for treating / diagnosing illness

59
Q

What is Heisenberg’s uncertainty principle?

A

It is impossible to know the exact position of an electron at a precise moment in time

60
Q

what provides evidence for sublevels?

A

ionisation energy

60
Q

what provides evidence for sublevels?

A

ionisation energy

61
Q

What type of structure does Si (Silicon) have?

A
  • metalloid
  • macromolecular covalent structure
  • very strong bonds
  • very high melting point
62
Q

define ionisation energy

A

energy required to remove one mole of electrons from 1 mole of gaseous atoms to produce 1 mole of gaseous 1+ ions

X (g) –> X+ (g) + e-

63
Q

Why are Be –> B and Mg –> Al exceptions to trend of ionisation energy?

A
  • B has higher NUCLEAR CHARGE than Be, but ionisation energy is lower
  • electrons in p-orbital are of HIGHER ENERGY & FURTHER FROM THE NUCLEUS than electrons in s orbital, hence they require SLIGHTLY less energy to remove
64
Q

Why are N –> O and P –> S exceptions to trend of ionisation energy?

A
  • S has lower ionisation than P
  • electron in same p-orbital as another electron is easier to remove than one in an orbital by itself due to REPULSION of other electron
65
Q

define electron affinity

A

energy RELEASED when one mole of electrons is added to one mole of gaseous atoms to produce one mole of gaseous 1- ions

X (g) + e- –> X- (g)

66
Q

is electron affinity exothermic or endothermic?

A

exothermic for FIRST electron, endothermic for SECOND electron

67
Q

How can alkali metals be prevented from reacting with air?

A

they can be stored under liquid paraffin

68
Q

How does alkali metals react w/ water?

A
  • lithium floats & reacts quietly
  • sodium melts into a ball which darts around on surface
  • potassium –> heat generation, ignites hydrogen
69
Q

why does melting points of ALKALI metals decrease down group 1 but increase down group 7?

A
  • WEAKER METALLIC FORCES FOR GROUP 1
  • STRONGER LDF FORCES FOR GROUP 7
70
Q

Which halogens for gaseous compounds with hydrogen?

A

HF, HI, HCl, & HBr (everything except Astatine)

71
Q

What do alkali metal halides react to form?

A

they are SOLUBLE & they react to form COLOURLESS, NEUTRAL SOLUTIONS

(it is the diatomic molecule that produces the colour)

72
Q

group 1 are …?

A

good REDUCING agents

73
Q

group 17 are . . .?

A

good OXIDIZING AGENTS

74
Q

does aluminium oxide react with wateR?

A

NO

75
Q

how can the presence of halide ions in solution be detected?

A

by adding silver nitrate solution - the silver ions react with the halide ions to form a precipitate of the silver halide –> silver halides can be distinguished by their color

76
Q

colours of silver halides

A

The silver chloride compound forms a white precipitate, silver bromide a creamy coloured precipitate and silver iodide a yellow coloured precipitate.

77
Q

what do metal oxides form on reacting with water?

A

metal hydroxides

78
Q

colors of the halogens?

A

Cl2 = green
Br2 = reddish
I2 = brown

Chemistry SL (5 decks)

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