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Periodicity Flashcards

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1
Q

What is the period number/ principle quantum number and what is its abbreviation

A

(n), the relative overall energy in each orbital

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2
Q

What are valence electrons

A

Electrons found in the outer shell

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3
Q

What is different about elements H and He

A

They dont comfortably fit into a specific group, so He is allocated group 0 based on shared similar physical and chemical properties. H doesn’t behave like other elements so is placed in its own group.

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4
Q

How do chemical and physical properties change across a period

A

They gradually change

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5
Q

How do chemical and physical properties act in the same group

A

Elements in the same group share similar physical and chemical properties

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6
Q

What are periodic trends

A

Specific patterns that are present in the periodic table that illustrate different aspects of a certain element

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7
Q

What are the four blocks all elements fall within

A

s, p, d ,f

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8
Q

S-block

A

All the elements with only s electrons in the outer shell

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9
Q

P-block

A

All the elements with at least 1 p electron in the outer shell

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10
Q

D-block

A

All the elements with at least 1 d electron and at least 1 s electron in outer shell, but no p of f (until 5d)

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11
Q

F-block

A

All the elements with one f electron and at least one s electron, but no d or p in outer shell

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12
Q

What is ionisation energy

A

The amount of energy required to remove an electron from an isolated ion/ molecule

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13
Q

First ionisation energy

A

The energy required to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions

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14
Q

What are the factors affecting ionisation energy

A

Atomic radius

Nuclear charge

Electron shielding

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15
Q

Atomic radius

A

Electrons in shells that are further away from the nucleus, have less attraction to it. Therefore, the further the outer shell is from the nucleus, the lower the ionisation energy. Because its easy to overcome these attractive forces when they are lower and take the electron

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16
Q

Nuclear charge

A

Nuclear charge increases with increasing atomic number. This means there are greater attractive forces between the nucleus and outer electrons. This means more energy is required to overcome the attractive forces when removing an electron so ionisation energy needs to be higher.

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17
Q

Electron shielding

A

The shielding effect is when electrons in the full inner shell, repel electrons in outer shells, preventing them from feeling the full nuclear charge. The greater the shielding, the lower the ionisation energy, because it’s easier to take an electron that feels little force keeping it to close to the inner shell.

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18
Q

What is the trend in FIE down the group

A

As you move down a group, FIE decreases

-atomic radius increases
-shielding by inner electrons increases
- therefore attraction between nucleus and outer electrons decreases
- so taking an e is easier

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19
Q

What is the trend in FIE across a period

A

FIE increases across a period, harder to take the e

-Nuclear charge increases across the period
- shielding remains the same
-atomic radius remains reasonably constant
- nuclear charge inc means its harder to take an e so ionisation energy needs to be higher

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20
Q

What happens to the FIE between the last element in one period and the first of another

A

Decreases

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21
Q

Why does FIE between the last element in one period and the first of another decrease

A
  • increased atomic distance- lower attractive forces- easier to take e
    -Increased shielding by inner electrons- outer electrons are repelled more so forces of attraction are also weaker- easier to take the e
  • these factors outweigh the increased nuclear charge
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22
Q

What is successive ionisation energy

A

When more and more electrons are removed, each from an ion that is becoming increasingly positive

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23
Q

What happens to the SIE of an element

A

Increases, as removing an electron from a positive ion is harder than removing one from a neutral atom.

-as more e are removed, attractive forces increase due to less shielding and an increased proton:electron number

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24
Q

What do big jumps on an IE graph show

A

Change of shell

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25
What do small jumps on an IE graph show
Change of sub shell
26
Characteristics of diamond (6)
-giant covalent lattice -macromolecule of C atoms -each c covalently bonded to 4 others -tetrahedral arrangement -bond angle is 109.5 -hardest known substance and therefore used in drills and glass cutting tools
27
Characteristics of graphite (6)
- each C bonded to covalently to 3 others -layers are hexagonal - bond angle of layers is 120 -spare electrons 1/4 are delocalised and occupy the space between layers -atoms in the same layer are held together by strong covalent bonds - layers held together by weak intermolecular forces which allow them to slide over eachother
28
Characteristics of graphene (3)
- infinite lattice of covalently bonded atoms in 2D to form only layers -single layer of C atoms bonded together in repeating patterns of hexagons - so thin it is considered 2D (1 million times thinner than paper)
29
Characteristics of silicon (IV) oxide (6)
-aka silicon dioxide - same structure as diamond, giant covalent lattice -macromolecule structure - tetrahedral shape bonded by strong covalent bonds -109.5 - each silicon shared by 4 oxygen, each oxygen shared by 2 silicons, SiO2
30
The 4 main physical properties
bp, mp, solubility, electrical conductivity
31
Mp and bp of metallic substances
High mp and bp lots of energy required to overcome the strong electrostatic forces of attraction between positive ions and a sea of delocalised electrons
32
Solubility of metallic substances
Metals dont dissolve. there are some interactions between polar solvents and charges in the metallic lattice but this can lead to reactions rather than dissolving
33
Electrical conductivity of metallic substances
Conduct electricity in both solid and and liquid states as delocalised electrons are free to move and carry charge
34
Mp and bp of giant covalent substances
Large number of covalent bonds holding together the whole structure, and lots of energy required to break it
35
Giant covalent lattices an be hard/soft, why
-graphite is soft due to wif between C layers -diamond and silicon dioxide are hard, because its very difficult to break the 3d network of strong covalent bonds -graphene is strong because of the strong covalent bonds, and it is flexible and used in transport
36
Solubility of giant covalent lattices
Most are insoluble in water
37
Electrical conductivity of giant covalent lattices
diamond and silicon dioxide dont conduct electricity and there are no free electrons available because alll 4 bonds of the carbon are filled. graphite and graphene do bc of the delocalised e being free to move and carry charge through the structure
38
What are group 2 elements commonly referred to as
Alkali metals, reducing agents
39
Why are group 2 elements considered reducing agents
-group 2 contains s2 configuration because it has 2 e in outer shell -redox reactions are common here -each metal atom is oxidised, losing its 2 valence e to form a 2+ ion with noble gas configuration -another species will gain the two e lost -that species is reduced
40
Explain the trends in reactivity of group 2 elements
- reactivity inc as you go down the group -bc as atomic radii increase there is more electron shielding - nuclear attraction then decreases which makes it easier to remove outer electrons -cations are more easily formed
41
Changes of oxidation number in the redox reactions
- becomes +2 when metal has been oxidised - becomes -2 when metal has been reduced
42
Reactions of group 2 metals with acids
react with dilute acids to form metal salts
43
If a group 2 is reacted with Hcl what would it form
metal chloride
44
What is the by product of the reaction between group 2 metals and acids
hydrogen gas
45
What is formed when group 2 metals react with dilute H2SO4 rather than Hcl
An insoluble sulfate
46
What happens down the group to group 2 sulphates
They become less soluble
47
Trends in IE of group 2
-the atoms of group 2 elements react by losing electrons to form 2+ ions - this formation of 2+ ions from gaseous atoms requires 2 ionisation energies -both first and second ionisation energies decrease down the group -as attraction between nuclear and outer electrons decreases due to inc atomic radius - inc atomic radius means more electron shielding
48
What happens when group 2 metals react with water
hydroxide ions are released alkaline solutions are formed
49
What happens to the solutions of group 2 metals when they react with water down the group and why
-they get more alkaline down the group -due to increasing solubility, meaning resulting solutions formed contain more hydroxide ions
50
How is a precipitate formed in group 2 metals
When g2 metals react with water they release hydroxide ions and form alkaline solutions. When the solution becomes even more saturated, the metal and OH- ions will from a precipitate
51
The two main uses of group 2 compounds
-agriculture -medicine
52
Group 2 compounds in agriculture, Calcium hydroxide (what is it used for, why is how much of it used important for regulation, why is It used)
-used as It is reasonably soluble in water -used to neutralise acidic soils -important to regulate how much Is added because if its too much the soil will become too alkaline to sustain crop growth
53
Group 2 compounds in medicine, Magnesium hydroxide (why's it used, used in what, why is it safe to use)
-used as its only partially soluble in water -used in suspension -to neutralise excess acid in the stomach and treat constipation -safe to use as only partially soluble, so only slightly alkaline due to the low OH- concentration
54
What type of molecules are halogens at room temp
Diatomic molecules
55
What type of structure are halogens and what forces do they have
- simple molecular structure -weak London forces between the diatomic molecules which are caused by instantaneous dipole- induced dipole forces
56
Process of instantaneous dipole- induced dipole
- electrons evenly distributed across a molecule, e.g I2 -an instantaneous dipole occurs when more of the e happen to be at one end more than the other - the instantaneous dipole in one I2 molecule induces a dipole in another -instantaneous dipole-induced dipole
57
Why do molecules become harder to separate, and bp increases as you down g7
- the more e there are in a molecule, the greater the instantaneous dipole- induced dipole forces -therefore the larger the molecule, the stronger the London forces and therefore more energy needed to overcome
58
Trend in reactivity of group 7
decreases as u go down
59
Cl, Br, I: Which potassium halides are added
potassium chloride potassium bromide potassium iodide
60
Colour of Cl when potassium chloride is added
Very pale green solution, often colourless
61
Colour of Br when potassium bromide is added
yellow solution
62
Colour of I when potassium iodide is added
brown solution, sometimes black solid is present
63
What colour is the organic solvent layer in the test tube of cl, br, I
cl: colourless br: yellow I: purple
64
Why does reactivity decrease down group 7
-oxidising power of the halogens decreases as you go down the group bc - atomic radius increases -more inner shells so shielding increases -less nuclear attraction to attract an electron from another species
65
Sub shell configuration of all halogens
2 electrons in outer s subshell, 5 in outer p sub shell
66
Why are halogens oxidising agents
- As they react with metals by accepting an electron to form a 1- ion - means oxidation number of the metal increases - halogens become reduced again as they gain an extra electron from the metal atom -oxidiation number of the halogen decreases
67
What is a disproportionation reaction
Reaction in which the same species is both oxidised and reduced
68
Examples of disproportionation reactions
-Cl with water - Cl with cold, dilute aq alkali
69
Chlorines uses in water
- clean water -make it drinkable
70
Chloric acid uses and uses of CLO-
- sterilises water by killing bacteria - can further dissociate in water to form CLO- -also acts as a sterilising agent
71
What happens if universal indicator is added to CLO-
-first turn red due to acidity of both reaction products -then turn colourless as HCLO bleaches the colour
72
Why is Cl used in water treatment to kill bacteria
-has been used to treat drinking water and swimming pool water -benefits outweigh risk of toxic. effects and possible risks from the formation of chlorinated hydrocarbons
73
How can halide ions be identified in an unknown solution
by dissolving the solution in nitric acid, then adding silver nitrate dropwise
74
Why is nitric acid aded in tested for halide ions
To prevent any false positive results from carbonate ions precipitating with the silver ones
75
Colour of precipitate for AgCl
White
76
Colour of precipitate for AgBr
Cream
77
Colour of precipitate for AgI
yellow
78
Adding ammonia is done why
To distinguish between the three colours as they can be very similar
79
How is ammonia added in halide ion testing
As dilute and then concentrated
80
results of adding dilute/ conc. ammonia
- If precipitate dissolves in dilute ammonia, unknown halide is chloride -If precipitate doesn't dissolve in dilute ammonia, but does in conc. its bromide -If precipitate doesn't dissolve in dilute or conc ammonia, its iodide
81
What is observed if carbonate Is present in a carbonate test and which gas is produced
effervescence, co2 produced and this forms a white precipitate of calcium carbonate when bubbled through limewater
82
What is calcium carbine bubbled through in the carbonate test
Limewater
83
What can be observed in a positive test for sulphates
White precipitate of barium sulfate is formed
84
What is used to do the test for ammonia ions
warm aq sodoium hydroxide
85
Whats formed in the positive test for ammonia ions
Ammonia gas
86
How can ammonia gas be distinguished
-pungent smell -turning red litmus paper blue

Chemistree (4 decks)

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