Chemistry (Tara) > Atomic theory 8 > Flashcards

Atomic theory 8 Flashcards

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1
Q

Define Atomic Radius

A

The atomic radius is obtained by getting half the distance between the centres of singly bonded atoms of the same element

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2
Q

Is there an increases or decrease in the atomic radius going across a period?

A

Decrease

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3
Q

Why is there a decrease in the atomic radius on going across a period?

A
  1. An increase in the nuclear charge (number of protons)

2. No change in the number of shells hence the screening effect of completed shells remains the same.

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4
Q

Is there an increase or decrease in the atomic radius on going down a group?

A

Increase

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5
Q

Why is there an increase in the atomic radius on going down a group?

A
  1. An increase in the number of shells
  2. The screening effect of electrons in the full inner shells blocks out some of the ‘pull’ that the nucleus has on outer electrons.
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6
Q

Define Electronegativity

A

The relative power of attraction an atom of an element has for the shared pair of electrons in a covalent bond.

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7
Q

What element has the greatest attraction for electrons in bond-forming situations?

A

Flourine

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8
Q

Can fluorine ever give electrons?

A

No its always gets them (unless made unnaturally)

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9
Q

Why does fluorine have the greatest attraction for electrons in bond-forming situations?

A
  1. The strongest positive nucleus

2. The smallest radius in its period = highest value

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10
Q

Does fluorine have the smallest or largest radius of all atoms?

A

Smallest therefore its best at attracting electrons

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11
Q

Which atom ‘pulls’ electrons away from which atom?(higher E or lower E?)

A

The atom with the higher electronegativity value will always ‘pull’ the electrons away from the atom that has the lower electronegativity value.

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12
Q

What does the degree of movement of the electrons toward the more electronegative atom depend on?

A

…The difference(subtraction) in electronegativities between the atoms involved

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13
Q

The bigger the difference in the electronegativity of the elements forming a bond, the more ……. the bond

A

Polar

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14
Q

E.D (greater or equal to) 1.7 =

A

Ionic

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15
Q

E.D < 1.7 =

A

Polar covalent

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16
Q

0 > E.D < 0.4 =

A

Slightly polar covalent

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17
Q

E.D = 0 =

A

Pure covalent

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18
Q

Does the electronegativity value generally increase or decrease going across a period?

A

Increase

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19
Q

Why does the electronegativity generally across a period?

A
  1. An increase in nuclear charge (number of protons increases)
  2. Leading to a decrease in atomic radius
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20
Q

Does the electronegativity value generally increase or decrease going down a group?

A

Decrease

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21
Q

Why does the electronegativity value generally decrease going down a group?

A
  1. The addition of extra shells making the atomic radius larger
  2. The screening effect of inner completed shells on outer electrons
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22
Q

Why do metals generally have low electronegativity values?

A

Because they are electropositive.

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23
Q

Does the atom with the lower electronegativity value usually get written first or last?

A

First eg. NaCL

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24
Q

What are cations?

A

When an atom loses electrons it becomes a positive ion (cation). The number of protons (+) remains the same but the number of electrons (-) decreases, hence we get an overall positive charge.

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25
What are anions?
When an atom gains electrons it becomes a negative ion (anion). The number of protons (+) remains the same but the number of electrons (-) increases, hence we get an overall negative charge.
26
Write in the correct symbol, valency and charge for the following: Mg ion K ion O ion
Mg(2+) K+ O(2-)
27
Define ionic bond
An ionic bond is formed due to the electrostatic attraction between oppositely charged ions caused by the complete transfer of electrons from one atom to another.
28
Does ionic bonding result in the formation of individual molecules?
No!
29
Do ionic molecules exist?
No!
30
Why are no molecules are present in ionic compounds?
Because no ion belongs specifically to any other ion but to all ions surrounding it!
31
What are ionic compounds made up of? (and why are they solid)
Made up of an orderly arrangement of oppositely charged ions held together by electrostatic forces.
32
Why are ionic compounds crystalline?
Because they are made up of an orderly arrangement of oppositely charged ions which give rise to a lattice structure.
33
Why are ionic compounds solid?
There are extremely strong attractions between the oppositely charged ions which hold the ions tightly together giving a solid structure.
34
Why do ionic compounds have high boiling points/melting points?
The extremely strong attractions between the oppositely charged ions require large amounts of energy to break, hence ionic compounds have high b.p and m.p.
35
What are 4 properties of ionic compounds (!!)
1 Usually solid crystalline 2. Conduct electricity when either in the molten or dissolved state as the ions are free to move. 3. (!!) Dissolve or are soluble in polar solvents, such as H₂O. 4. High boiling points and melting points.
36
Why do ionic compounds dissolve in water?
As the attraction between the ions in the compound and polar water are strong enough to pull the crystalline lattice apart.
37
Name 3 samples of ionic materials in everyday life.
1. Washing soda crystals (Na2CO3) = soften hard water e.g. Calgon tablets 2. Aluminium sulphate [Al2(S04)3] = flocculating agent in the preparation of water for drinking. 3. Salt tablets are used to replace salt lost by sweat.
38
What shape is water?
V-shaped
39
What leads to a molecule being polar?
Lack of symmetry
40
What does each kind of atom have?
A certain attraction for the electrons involved in a chemical bond.
41
Why are all other atoms given values less than that of fluorine?
As fluorine has the smallest radius of all the aomts.
42
What happens when atoms react with each other?
They compete for the electrons in the bond.
43
Does fluorine ever give electrons?
No it always gets them (unless made unnaturally)
44
With a small radius and a relatively large nuclear charge the atom has …….
Great attraction for electrons
45
With a large radius the nucleus finds it ……..
More difficult to attract electrons.
46
Why does the O in OF₂ come first but come last in Cl₂O?
As the atom with the lower electronegativity value usually gets written first.
47
Intramolecular means
'within one molecule'
48
Intermolecular means
'between two or more molecules'
49
What do simple anions names end in?
'ide'
50
What does the arrangement of ionic compounds give rise to?
This arrangement give rise to a crystalline structure called a LATTICE.
51
What is a lattice?
A crystalline structure
52
Does an ion belong specifically to any other ion?
No but to all other ions surrounding it.
53
Are molecules present in ionic compounds?
No
54
Where does each positive ion exert its positive charge?
In all directions around itself and so attracts many of the negative ions to itself.
55
Where does each negative ion exert its negative charge?
In all directions around itself and so attracts many of the positive ions to itself.
56
Can the words ionic and molecules be used in the same sentences?
NO
57
Tightly packed =
Solid
58
Why are ionic compounds usually solid crystalline?
Due to the individual ions attracting other ions (extremely strong forces of attraction)
59
Why do ionic compounds conduct electricity when either in molten or dissolved state?
As the ions are free to move.
60
Why do ionic compounds dissolve or are soluble in polar solvents, such as H₂O?
As the attraction between the ions in the compound and polar water are strong enough to pull the crystalline lattice apart.
61
Why do ionic compounds have high melting a nd boiling points?
Due to a lot of energy being needed to break the extremely strong forces of attraction between the ions.
62
Give a sample of a flocculating agent.
Aluminium sulphate, Al₂(SO₄)₃, flocculating agent in the preparation of water for drinking. It causes small suspended particles to stick together so that they can be filtered off easer.
63
Name a disease that aluminium is associated with.
Alzheimer's
64
Define electronegativity (L.C)
The relative power of attraction an atom of an element has for the shared pair of electrons in a covalent bond.
65
Why is there an increase in electronegativity value moving from gallium to germanium in the periodic table? (L.C)
- Nuclear charge increasing | - Atomic radius decreasing
66
Mendeleev predicted the properties of the elements gallium and germanium years before either of them discovered. Explain the basis for his predictions. (L.C)
Predicted the properties of known elements
67
Write the molecular formula for the simplest compound formed between germanium and hydrogen. (L.C)
GeH₄
68
Would you expect GeH₄ to be water soluble? Justify your answer. (L.C)
- No | - GeH₄ is a non-polar solute
69
Define electronegativity (L.C)
The relative power of attraction an atom of an element has for the shared pair of electrons in a covalent bond.
70
State and explain the trend in electronegativity values down the first group in the periodic table of elements. (L.C)
- Decrease | - Increasing atomic radius and increased screening offsets increased nuclear charge
71
Use electronegativity values to predict the types of bonding (i) in water, (ii) in methane, (iii) in magnesium chloride. (L.C)
Water - polar covalent Methane - covalent Magnesium chloride - ionic
72
Use your knowledge of intermolecular forces to explain why methane has a very low boiling point. (L.C)
Very weak intermolecular forces.
73
The relative molecular mass of methane is only slightly lower than that of water but the boiling point of water is much higher. Suggest a reason for this (L.C)
Much stronger hydrogen bonds between water molecules.
Chemistry (Tara)
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