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Chemistry > Topic 3: The periodic table > Flashcards

Topic 3: The periodic table Flashcards

1
Q

Which way are groups and which was are periods?

A

Groups are the vertical columns
Periods are the horizontal rows

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2
Q

Roughly where are metals, non metals and metalloids on the periodic table? (imagine)

A

Yellow: Metals
Pink: Non metals
Green: Metaloids

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3
Q

What are the elements called when they have both properties of both metals and non metals?

A

Metalloids

shaded green are metalloids
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4
Q

What are metalloids?

A

Elements when they have both properties of both metals and non metals

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5
Q

Which is the most abundant element in the Universe: about 90% of the atoms?

A

Hydrogen

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6
Q

What are noble gases also known as?

A

Inert gases

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7
Q

What are the elements that have been omitted called?

A

Lanthanoid elements and actinoid elements

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8
Q

What does a metallic structure consist of?

A

consists of a regular lattice of positive ions in a sea of delocalised electrons

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9
Q

What can the metallic and non-metallic properties of elements be related to? What energy?

A

Ionisation energies

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10
Q

What must the element do to form a metallic structure?

A

They must be able to lose electrons fairly readily to form positive ions

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11
Q

How does the ionisation energy change across a period and down a group?

A

Across: ionisation energy increases → so elements lose electrons less easily

  • so metallic structures are formed by elements on the left hand side of the periodic table, which have lower ionisation energies

Down: Ionisation energy decrease → therefore elements are much more likely to exhibit metallic behaviour lower down a group

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12
Q

What are the properties a metallic element would generally have?

A
  • Large atomic radii
  • Low ionisation energies
  • Less exothermic electron affinity values
  • Low electronegativity
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13
Q

How are valence electrons presented on the periodic table?

A

As groups

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14
Q

For groups 13-18, how are valence electrons shown?

A

group number - 10

ex. elements in group 13 have 3 valence electrons

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15
Q

What does the period number indicate?

A

The number of shells (main energy levels) in the atoms - or which is the outer shell (main energy level)

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16
Q

What is the periodic table divided into blocks according to?

A

According to the highest energy sub-shell (sub-level) occupied by electrons.

→ in the s block, all elements have atoms in which the outer shell electron config. is ns1 or ns2 (n is the shell number) and in the p block it is the p subshell that is being filled

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17
Q

What are the physical properties that can be deduced from the periodic table?

A
  • Atomic radius
  • Ionic radius
  • First ionisation energy
  • Electron affinity
  • Electronegativity
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18
Q

What does the atomic radius describe?

A

The size of an atom

The larger the atomic radius, the larger the atom

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19
Q

How does atomic radius change across and down a period?

A

Across: Decreases across a period, because nuclear charge increases across the period with no significant change in shielding → Meaning that the outer electrons are pulled in more strongly and the radius is smaller

Down: Increase because atoms have increasingly more shells

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20
Q

What does the ionic radius measure?

A

The size of an ion

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21
Q

In general, are positive/negative ions bigger or smaller than their atomic radii?

A

In general, the ionic radii of positive ions are smaller than their atomic radii, and the ionic radii of negative ions are grater than their atomic radii

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22
Q

Why does Na have a bigger ionic radius than Na+?

A

Because Na has one extra shell whereas Na+ lost one.

There is also a greater amount of electron-electron repulsion in Na because there are 11 electrons compared to 10. The electron cloud is therefore larger in Na than in Na+ because there are more electrons repelling for the same nuclear charge pulling the electrons in.

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23
Q

How does ionic radius change across a period?

A

It is not a clear-cut trend because the type of ion changes going from one side to the other

Positive ions are formed on the left hand side and the negative ion on the right hand side.

Even though on the same period (have same shells), there are just more valence electrons on the right hand side (negative) for there to be a stronger election-electron repulsion

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24
Q

How does the ionic radius change for negative and positive ions?

A

For positive ions, there is a decrease in ionic radius as the charge on the ion increases

For negative ions, the size increases as the charge increases

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25
Which one has a larger ionic radius? Na+ or Mg2+?
Both ions have the same electron configuration, same number of electrons → same electron-electron repulsion However there is one more proton in Mg2+ and the higher nuclear charge in Mg2+ means that the electrons are pulled in more strongly and so the ionic radius is smaller
26
What is the definition of first ionisation energy?
The energy required to remove the outermost electron from a gaseous atom
27
How does the first ionisation energy change down and across a period?
Down: the first ionisation energy decreases Across: the general trend is that it increases from left to right due to increase in nuclear charge across the period
28
Why is the first ionisation energy decreasing down a group?
Because there are more shells, size of the atom increase so the outer electron is further from the nucleus and therefore less strongly attracted by the nucleus Although the nuclear charge also increases, it is largely balanced out by the increase in shielding down the group.
29
What governs the change in first ionisation energy?
the increase in size
30
What is the general trend for first ionisation energy from left to right across a period?
The general trend is that first ionisation energy increases from left to right across a period. This is because of an increase in nuclear charge across the period (protons are added)
31
Why does the first ionisation energy increase across a period?
the nuclear charge increases as protons are added to the nucleus. Therefore the attractive force on the outer electrons due to the nucleus increases from left to right across the period.
32
Explain the exception to the general increase in first ionisation energy across a period
The major difference is that the electron to be removed from the boron atom is **2p sub-level** whereas it is in a **2s sub-level** in beryllium. The 2p sub-level in B is higher in energy than 2s sub-level in beryllium, therefore less energy is required to remove an electron from boron
33
What are the two exceptions in first ionisation energies?
Boron vs beryllium * Main difference is that the electron to be removed in born is in the 2p sub-level, whereas it is in 2s sub-level in beryllium. The sp sub-level in B is higher in energy than the 2s sub-level in beryllium, therefore requiring less energy to remove an electron from boron Oxygen vs nitrogen * the first ionisation energy of oxygen is lower than that of nitrogen * The major difference is that oxygen has two electrons paired up in the same p orbital, but nitrogen does not. An electron in the same p orbital as another electron is easier to remove than one in an orbital by itself because of the repulsion from the other electron. * When two electrons are in the same p orbital they are closer together than if there is one in each p orbital. If the electrons are closer together, they repel each other more strongly. If there is a greater repulsion, an electron is easier to move
34
What is electron affinity defined as?
It involves the energy change when one electron is added to a gaseous atom the enthalpy change when one electron is added to each atom in one mole of gaseous atoms under standard conditions
35
Why is the first electron affinity is exothermic for virtually all elements?
It is a favourable process to bring an electron from far away to the outer shell of an atom, where it feels the attractive force of the nucleus
36
How does the electron affinity change down a group and across a period?
Down: electron affinity decrease, becomes less exothermic as the size of the atom increases. There is a weaker attraction between the added electron and the nucleus if the atom is big (brought to a position which is further from the nucleus) Across: the electron affinity becomes more exothermic. This is because of an increase in nuclear charge and a decrease in atomic radius from left to right across the period. → electrons will be more strongly attracted when brought into the outer shell of the atom
37
Which element has the most exothermic value for electron affinity?
Chlorine
38
What is electronegativity a measure of?
the attraction of an atom in a molecule for the electron pair in the covalent bond of which it is a part
39
What is how strongly the electrons are attracted dependent on?
the size of the individual atoms and their nuclear charge
40
How does the electronegativity change down a group and across a period?
Down: Decrease because the size of the atom increases down a group Across: Increases because there is an increase in nuclear charge across the period with no significant change in shielding
41
What are the elements in group 1 known as?
Alkali metals
42
What are the properties of alkali metals?
* Highly reactive * Soft * Low melting point metals
43
What are the 2 reasons alkali metals are placed together in group 1?
* They all have one electron in their outer shell * They react in very similar ways (similar chemical properties)
44
What is the reaction of an element determined by?
The number of electrons in the outer shell of their atoms. → elements in the same group in the periodic table have the same number of electrons in their outer shell, they react in basically the same way
45
What is the bonding in group 1 elements and how are the solid held together?
* All metallic bonding * Solid is held together by electrostatic attraction between the positive ions in the lattice and the delocalised electrons
46
What is electrostatic attraction?
When negatively charged atom is attracted towards positively charged atom and vice-versa, it is known as electrostatic attraction.
47
What is the electrostatic attraction due to?
The attraction for the delocalised, negatively-charged, electrons is due to the nucleus of the positive ion.
48
What happens to the ions in group 1 elements as they go down the group?
* ions get larger * nucleus becomes further from the delocalised electrons and the attraction becomes weaker * → less energy is required to break apart the lattice going down group 1
49
What is the trend for melting point down group 1?
It decreases
50
Why do reactions become more vigorous going down the group?
Because the ionisation energy decreases as the size of the atom increases. E.g. caesium oses its outer electron to form a positive ion much more easily than sodium and will react more vigorously.
51
How does group 1 elements react with oxygen and water? How does the reaction with water change as it goes down a group
Oxygen: The alkali metals react vigorously and all tarnish rapidly in air Water: React rapidly. Alkaline solution is formed. The alkali metal hydroxides are strong bases and ionise completely in aqueous solution Reaction with water becomes **more vigorous** going **down the group**
52
What are the elements in group 17 known as?
Halogens - all non metals consisting of diatomic molecules (X2)
53
Properties of halogens (F, Cl, Br, I) M/B point/ colour/ physical state at room temp. and pressure
54
How does the melting points of halogen change down the group?
It increases because london forces between molecules get stronger → more energy must be supplied to separate the molecules from each other
55
How does the reactivity of halogens change down the group?
Down: decreases
56
What is the most reactive element in group 17?
Fluorine (halogen)
57
How can the very high reactivity of fluorine be explained?
Can be explained in terms of an exceptionally weak F-F bond and the strength of the bonds it forms with other atoms. The reactivity in terms of the formation of X- ions can be related to a decrease in electron affinity (energy released when an electron is added to a neutral atom) going down the group as the electron is added to a shell further away from the nucleus
58
What do halogens from when reacting with alkali metals?
Salts * All white/colourless * fairly typical ionic compounds
59
What do all alkali metal chlorides, bromides and iodides have in common and form?
* Soluble in water * Form colourless, neutral solutions
60
What does the vigorousness of the reaction depend on for halogens?
How vigorous the reaction is depends on the particular halogen and alkali metal used
61
What are displacement reactions of halogens?
These are reactions between a solution of a halogen and a solution containing halide ions A small amount of a solution of a halogen is added to a small amount of a solution containing a halide ion, and any colour changes are observed
62
What will the results of reactions between halogen solutions and solutions containing halide ions be? (Cl2, Br2, I2 vs KCl, KBr, KI)
63
Why would a solution turn orange?
Bromine solution - Due to the production of bromine
64
Why would a solution turn red-brown
Iodine solution - Due to the production of iodine
65
What is the colour of a chlorine solution?
Pale yellow-green
66
In a reaction, what determines which halogen displaces the halide ion of another halogen?
The more reactive halogen displaces the halide ion of the less reactive halogen from the solution. e.g. chlorine displaces bromide ions and iodide ions from solution, and bromine displaces iodide ions from solution → Cl is a stronger oxidising agent than bromine and iodine so it will oxidise bromide ions to bromine and iodide ions to iodine → Bromine is a stronger oxidising agent than iodine and will oxidise iodide ions to iodine → in terms of electrons, Cl has the strongest affinity for electrons and will remove electrons from bromide ions and iodide ions
67
What does a basic oxide reacted with an acid form?
Salt, and if soluble in water, will produce an alkaline solution
68
What can oxides of elements classified as? (3)
basic, acidic or amphoteric
69
Are these basic or acidic? Metallic oxide vs non-metallic oxides
Metallic oxides: basic Non-metallic oxides: acidic
70
What element has properties of a basic oxide and some of an acidic oxide?
Aluminium is exhibiting properties between those of a metal (basic) oxide and those of a non-metal (acidic) oxide
71
What is a basic oxide?
is one that reacts with alkalis to form a salt and, if soluble in water, will produce an acidic solution
72
What does sodium oxide reacted with water form?
Sodium hydroxide
73
What do sodium oxides react with to form salts?
Acids such as sulfuric acid
74
Why is magnesium oxide not very soluble in water?
Because of the relatively high charges on the ions, but it does react to a small extent to form a solution of magnesium hydroxide, which is alkaline
75
How is Aluminium on the dividing line between metals and non-metals and forms an amphoteric oxide?
They have some of the properties of a basic oxide and some of an acidic oxide. Aluminium is exhibiting properties between those of a metal (basic) oxide and those of a non-metal (acidic) oxide
76
How does aluminium oxide display amphoteric behaviour?
It does not react with water but it does display amphoteric behaviour in that is reacts with both acids and bases to form salts
77
What do amphoteric oxides react with?
Both acids and bases
78
What is an acidic oxide?
An acidic oxide is one that reacts with bases/alkalis to form a salt and, if soluble in water, will produce an acidic solution
79
What does P4O6 (phosphorus (III) oxide) and P4O10 (phosphorus (V) oxide) form when reacting with water?
Phosphoric (III) and Phosphoric (V) acid
80
What does SO2 (sulfuric (IV) oxide) and SO3 (sulfur (VI) oxide) form when reacting with water?
Sulfuric (IV) and Sulfuric (VI) acid
81
What are the two most environmentally important nitrogen oxides?
Nitrogen (II) oxide (NO) and nitrogen (IV) oxide (NO2)
82
What is formed when nitrogen reacts with oxygen at very high temperatures?
to form NO (Nitrogen monoxide, nitric oxide or nitrogen (II) oxide)
83
Where does this reaction occur?
In the internal combustion engine
84
Is NO soluble and what is it classified as?
NO is virtually insoluble in water and is classified as a neutral oxide
85
What can NO be oxides into? Where does this happen?
NO can be oxidised in the atmosphere to NO2
86
What can NO2 form when reacting with water?
forms nitric (V) acid (HNO3), which is one of the acids responsible for acid deposition.
87
What can nitric oxide be classified as?
An acidic oxide
88
What is N2O?
(nitrogen (I) oxide, nitrous oxide) is another neutral oxide Also known as laughing gas and major uses include as an anaesthetic and as the propellant in ‘squirty cream'

Chemistry (12 decks)

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