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Chemistry > Module 3 > Flashcards

Module 3 Flashcards

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1
Q

define first ionisation energy

A

the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms

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2
Q

what factors effect an elements ionisation energy

A

atomic radius, nuclear charge, sheilding

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3
Q

define second ionisation energy

A

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

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4
Q

trends in ionsiation energy across a period

A

increases going from left to right. Atomic radius decreases, nuclear charge increases but shielding stays the same so there are stronger electrostatic forces.

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5
Q

trends in ionisation energy down a group

A

decreases as more electron shells so increased shielding and increased atomic radius causing weaker electrostatic forces. This counteracts the increase in nuclear charge.

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6
Q

Trend in MP/BP across period 2 and 3.

A

increases from group 1-14. (stronger metallic bonds as more positive ions are formed or giant covalent structures). Sharp decrease from group 14 - 15 (change from giant to simple structures). comparatively low from 15-18.

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7
Q

How can you tell how many electrons an element has in its outer shell

A

Looking at the ionisation energies. The biggest jump in ionisation energy is where it loses an electron shell e.g. if there is the biggest jump between 4 and 5 ionisation energy then there were 4 electrons in the outer shell.

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8
Q

why does ionisation energy drastically increases once all the electrons in the outer shell have been removed

A

decrease in shielding, decrease in atomic radius and increase in effectiveness of nuclear charge.

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9
Q

trend in reactiveness of group 2 elements

A

as you go down the group it is more reactive as it is easier to lose electrons

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10
Q

write a general equation for the reaction of water with a group 2 metal using M to represent the metal

A

M(s)+ 2H2O(l) —– M(OH)2(aq) + H2(g)

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11
Q

write an equation to show formation of metal hydroxide from a group 2 metal oxide and water

A

CaO(s) + H2O(l) —– Ca(OH)2(s)

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12
Q

solubility of group 2 metal hydroxides trend

A

as you go down the group solubility/alkalinity increases

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13
Q

chemical formula for lime water

A

Ca(OH)2

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14
Q

what is slaked lime and what is it used for

A

calcium hydroxide - added to fields of lime to increase pH of acidic soils

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15
Q

what calcium compound is used in indigestion treatments

A

calcium carbonate

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16
Q

describe test for sulphate ions

A

Acidify BaCl2 by adding a few drops of HCl. Add this to the solution. If sulphate ions are present then a white precipitate of BaSO4 forms.

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17
Q

state the chemical name and the use for barium meal. Why can barium not be swallowed but barium meal can

A

barium sulfate - detects abnormalities in the esophagus, stomach and small bowel. Bariums sulfate is not toxic unlike barium.

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18
Q

chemical formula and use for ‘milk of magnesia’ and why it has this name.

A

Mg(OH)2 - treats acid indigestion. Has a milky white colour.

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19
Q

trend in reactivity of group 17

A

reactivity decreases going down the group (harder to gain an electron)

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20
Q

boiling point trend of group 17

A

increases going down the group as there are stronger london forces as there are more electrons.

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21
Q

what is a halide

A

any group 17 element that has formed an ion

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22
Q

ionic equation for test for halides

A

Ag+(aq) + X-(aq) —- AgX(s)

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23
Q

describe the test for halides

A

add aqueous silver nitrate to aqueous solution. If halide present precipitate forms. Add aqueous ammonia to test solubility of precipitate

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24
Q

colour of precipitate and solubility of Cl in NH3

A

white. Soluble in dilute NH3

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25
Q

colour of precipitate and solubility of Br in NH3

A

cream. Soluble in concentrated NH3

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26
Q

colour of precipitate and solubility of I in NH3

A

yellow. Insoluble in concentrated NH3

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27
Q

ionic equation for carbonate test

A

2H+(aq) +CO3(2-)(aq) —— CO2(g) + H2O(l)

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28
Q

describe the carbonate test

A

add dilute nitric acid to solution. If you see bubbles it could be carbonate. Bubble gas through lime water if it turns cloudy then carbonate present.

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29
Q

ionic equation for sulfate test

A

Ba2+(aq) + SO4(2-)(aq) —— BaSO4 (s)

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30
Q

describe the test for sulfate ions

A

add barium nitrate to solution. If white precipitate forms sulfate ions are present

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31
Q

what are the sequence of tests

A
  1. carbonate, 2.sulfate, 3.halides
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32
Q

What is the order of the anion tests:

A

carbonate test is first as effervescence is observed unlike any of the others so no possible incorrect conclusion. Sulfate test second because a white precipitate also occur for carbonate and its already ruled out. Halide last and as precipitates also form so it would have to be a halide.

33
Q

how would you carry out the tests if there is a mixture of ions

A

carry put tests for same sequence. for carbonate test add nitric acid till all bubbling stops (all carbonate ions removed). For sulfate test add excess Ba(NO3)2 and filter off precipitate. Halide test as normal

34
Q

equation for ammonium ion test

A

NH4+(aq) + OH-(aq) —– NH3(g) + H2O(l)

35
Q

describe ammonium ion test

A

NaOH is added to NH4+. NH3 gas is produced which is alkaline so using indicator paper it will turn blue

36
Q

what is a disproportionation reaction

A

oxidation and reduction of the same element

37
Q

example of disproportionation reaction with NaOH and Cl

A

Cl2(g) + 2NaOH(aq) —- NaClO(aq) + NaCl(aq) + H2O(l). NaOH must be cold and dilute.

38
Q

example of disproportionation reaction with H2O and Cl

A

Cl2(g) + H2O(l) —- HClO(aq) + HCl (aq)

39
Q

what is the use for chloric (I) acid

A

bactericidal so can disinfect or purify water

40
Q

what are the risks associated with chlorinating water

A

chlorine can react with organic matter to form chlorinated hydrocarbons that can cause cancer. chlorine is a respiratory irritant

41
Q

what is NaClO

A

bleach

42
Q

energy change in the formation and breaking of bonds

A

energy supplied to break bonds. Energy released to form bonds

43
Q

exothermic reactions

A

give out heat energy. Products have less energy than the reactants

44
Q

endothermic reactions

A

takes heat. Products have more energy than the reactants

45
Q

define activation energy

A

minimum energy required for a reaction to take place

46
Q

energy change formula

A

Q=mc(change in)t

47
Q

what is standard enthalpy change of a reaction

A

enthalpy change when the amount shown in the chemical equation react under standard conditions with the reactants and products in their standard states

48
Q

what is standard enthalpy change of neutralization

A

the enthalpy change of reaction when amounts of an acid and an alkali react under standard conditions

49
Q

what is the standard enthalpy change of combustion of a substance

A

enthalpy change when one mole of the substance completely burns in oxygen under standard conditions with the reactants and products in their standard solution

50
Q

what is the standard enthalpy change of formation of a compound

A

enthalpy change when one mole of the compound forms from its elements under standard conditions with the elements and the compound in their standard states

51
Q

what is the bond dissociation enthalpy

A

enthalpy change on breaking one mole of a particular covalent bond in a gaseous mole.

52
Q

Describe the determination of an enthalpy change of combustion practical

A
  1. measure some water in a measuring cylinder. Pour into beaker and record initial temperature.
  2. Add methanol to spirit burner, weigh spirit burner containing ethanol.
  3. Place spirit burner under beaker. Light the burner and burn the methanol while stirring the water wth the thermometer.
  4. After 3 minutes extinguish flame and record max temperature reached.
  5. Reweigh the spirit burner
53
Q

energy change of combustion formula

A

energy transferred to water / mol

54
Q

enthalpy of combustion practical sources of error

A

heat loss to surroundings, incomplete combustion, evaporation of methanol, non-standard conditions (all lead to lower than expected results)

55
Q

limitations of enthalpy change of neutralization experiment

A

heat loss to surroundings, under non-standard conditions and specific heat capacity of cup not taken into account.

56
Q

what is the first law of thermodynamics

A

energy cannot be created or destroyed, only changed from one form to another

57
Q

why are enthalpy changes measured under standard conditions and what are the standard conditions

A

data is compared so conditions need to be standardised to ensure nothing else is effecting the enthalpy changes. 298K. 100KPa

58
Q

what is first electron affinity?

A

the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions.

59
Q

What is hess’ law

A

enthalpy change is independent of the route taken, providing the conditions are constant. If all but one enthalpy change is known then an expression can be formed to work out the missing enthalpy change

60
Q

what factors affect rate

A

concentration, pressure, temperature, catalyst and surface area

61
Q

what does increase concentration/pressure do to rate?

A

more particles in a given volume leads to higher frequency of successful collisions

62
Q

what does increase temperature do to rate?

A

increase KE/ more particles have necessary activation energy therefore higher frequency of successful collisions

63
Q

how can you measure the rate of reaction

A

volume of gas produced, mass of reactant and colour change

64
Q

what is homogenous catalyst

A

catalyst in the same physical state as the reactants

65
Q

what is heterogenous catalyst

A

catalyst in different physical state from the reactants

66
Q

how does a catalyst speed up the rate of reaction

A

lowers activation energy by finding an alternative reaction pathway

67
Q

what does the total area under the boltzmann distribution graph

A

total number of molecules

68
Q

what is dynamic equilibrium

A

stage in a reversible reaction where the rate of the forward reaction is equal to the rate of the backwards reaction. Concentration of reactants/products stay the same

69
Q

what is le chatelier’s principle

A

when a change is made to favour forwards or backwards reaction the system readjusts it to minimize the effects of that change.

70
Q

how do we know a reaction hasn’t stopped during equilibrium

A

by using radioisotopes

71
Q

what does equilibrium constant show

A

whether the reaction is more product or more reactant favored

72
Q

what is the Kc formula for this reaction A+B —- 2C

A

[C]2 / [A][B]

73
Q

what does it mean of Kc is greater than one

A

product favoured reaction

74
Q

what does it mean of Kc is less than one

A

reactant favoured reaction

75
Q

what can effect Kc

A

change in temperature

76
Q

what are the labels for the axis on a boltzmann distribution graph

A

y axis - number of molecules with a given energy

x axis - energy

77
Q

How does the boltzmann distribution graph depending on temperature. Why does this change occur

A

At higher temperature the peak is shifted to the right and the peak isn’t as high on the y axis. When it dips down the curve doesn’t dip as low as the lower temperature curve. More particles have activation energy so higher frequency of successful collisions.

78
Q

How do you prepare a standard solution

A
  1. weigh the solid on a weighing boat.
  2. Solid is dissolved in a beaker using less water than is needed to fill the volumetric flask.
  3. Solution is transferred to a volumetric flask. Last traces of the solution are rinsed into the flask with distilled water.
  4. The flask is carefully filled to the graduation line by adding distilled water a drop at a time until the bottom of the meniscus lines up exactly with the mark.
  5. Volumetric flask is inverted several times to mix the contents.

Chemistry (24 decks)

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