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chemistry > module 3 > Flashcards

module 3 Flashcards

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1
Q

define the first ionisation energy

A

energy required to remove one electron from each atom in one mole of gaseous atoms of an element

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2
Q

what electron will experience least nuclear attraction

A

highest energy level

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3
Q

how does atomic radius affect the ionisation energy

A

F of attraction falls off sharply with increasing distance

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4
Q

how does nuclear charge affect the ionisation energy

A

more protons in the nucleus, the greater the nuclear attraction of the outermost electron

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5
Q

how does electron shielding affect the ionisation energy

A

inner-shell e- repel outer-shell e-(shielding effect) which reduces the nuclear attraction of the outer electrons

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6
Q

how does an unpaired electron increase the ionisation energy

A

single electron pulled closer to the nucleus so the nuclear attraction increases so the ionisation energy increases

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7
Q

what is the trend in ionisation energy across the period

A

increase

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8
Q

what is the trend in ionisation energy for periods at the end of the table and periods at the start

A

sharp decrease

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9
Q

what is the trend in ionisation energy down the group

A

decrease

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10
Q

explain the trend in first IE down the group

A

atomic radius increases

shielding increases

outweighs the effect of the nuclear charge
so nuclear attraction decreases

so first IE decreases

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11
Q

explain the trend in first IE across a period

A

nuclear charge increases

(similar shielding)

nuclear attraction increases

atomic radius decreases

first IE increases

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12
Q

explain the lower IE of oxygen compared to nitrogen

A

outer electrons in same sub-shell

oxygen has a paired e- , the e- repel each other making it easier to remove

therefore, the IE of oxygen is less

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13
Q

outline the properties of metals

A

strong metallic bonds

high electrical conductivity

high melting & boiling point

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14
Q

what does the melting point of a metal depend upon

A

the strength of the metallic bonds

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15
Q

do metals dissolve

A

no

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16
Q

outline the melting&boiling points, the solubility and electrical conductivity of giant covalent structures

A

high melting & boiling points

insoluble

non-conductor of electricity (except graphene)

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17
Q

describe the trend in melting point across period 2 and 3

A

increases from group 1-14

sharp decrease between 14-15

m.p. remains low from 15-18

18
Q

define a reducing agent

A

something that reduces another species

19
Q

group 2 elements + oxygen –>

A

metal oxide

20
Q

group 2 elements + water –>

A

alkaline hydroxide + hydrogen gas

21
Q

explain the increase in reactivity down group 2

A

nuclear attraction decreases

as a result of more atomic radius and more shielding

22
Q

when the alkaline hydroxide produced becomes saturated, what happens

A

M2+ + 2OH- –> M(OH)2 (s)

precipitate forms

23
Q

what is the use of calcium hydroxide

A

neutralises acidic soils

24
Q

what are group 2 hydroxides and carbonates used for

A

digestion tablets

neutralises stomach acids

25
Q

what happens when a solution of chlorine is added to an aqueous solution of bromide and iodide ions

A

chlorine is more reactive

the halogen displaces the halide from solution

changing the colour of the solution

26
Q

what do chlorine, bromine and iodine make the solution in water

A

chlorine – pale green

bromine – orange

iodine – brown

27
Q

what do chlorine, bromine and iodine make the solution in cyclohexane

A

chlorine – pale green

bromine – orange

iodine – violet

28
Q

a solution of chlorine is added to a solution containing iodide ions

what’s the colour?

A

purple

iodine is formed
(chlorine displaces)

29
Q

Cl2 + H2O –>

A

HClO + HCl

30
Q

what causes the indicator paper to bleach after it turns red when added to a solution of water and chlorine

A

chloric(I) acid acts as a weak bleach

31
Q

how does chlorine kill bacteria

A

chlorine and water form HClO and HCl

bacteria are killed by the HClO and ClO-

32
Q

Cl2 + 2NaOH –>

A

NaClO + NaCl + H2O

33
Q

compare the reaction with water and chlorine, and chlorine and NaOH

A

ClO- formed at a larger concentration when chlorine reacts with NaOH

34
Q

describe the test for a carbonate

A

add d. nitric acid

bubble the gas produced through lime water (aqueous Ca(OH)2)

lime water turns milky (calcium carbonate produced)

35
Q

describe the test for sulfate

A

add barium ions to solution forming a ppt.

Ba2+ + SO4 2- –> BaSO4

36
Q

why is carbonate test performed first

A

neither sulfate nor halide ions produce bubbles with dilute acid

37
Q

why is the sulfate test done after the carbonate test

A

BaCO3 is also white

38
Q

why is the halide test performed last

A

Ag2CO3 and Ag2SO4 form ppt

39
Q

when testing for a mixture of ions, how do you tweak each test

A

carbonate test
continue adding d. nitric acid until the bubbling stops

sulfate test
add excess of Ba(NO3)2

halide test
add AgNO3 before doing the test, then add NH3 to confirm which halide

40
Q

describe the test for ammonium ions

A

add NaOH

(ammonia gas produces)

warm solution to release ammonia gas

test with moist indicator paper

chemistry (6 decks)

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