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Higher Chemistry > UNIT 1 > Flashcards

UNIT 1 Flashcards

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1
Q

What does a higher density mean?

A

The more closely packed the particles are

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2
Q

Why do gaseous elements have lower densities?

A

Because their particles are packed far apart

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3
Q

Why do solid elements have higher densities?

A

Because their particles are packed closely together

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4
Q

What does it mean for the bonds if the particles are packed together?

A

The bonds between the particles are likely to be stronger

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5
Q

What happens to the density of metals going across a period?

A

The density increases

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6
Q

What happens to the density of metals going down a group?

A

The density increases

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7
Q

What happens to the density of non metals going across a period?

A

There is no real pattern

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8
Q

What happens to the density of non metals going down a group?

A

The density increases

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9
Q

What is metallic bonding?

A

Metallic bonding is the strong electrostatic force of attraction between the positive charged ions

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10
Q

What are intramolecular forces?

A

The forces that hold atoms together within a molecule

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11
Q

What are intermolecular forces?

A

The forces that exist between molecules

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12
Q

What is covalent radius?

A

A measure of the size of the atom

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13
Q

What happens to covalent radius going across a period?

A

The covalent radius decreases

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14
Q

Why does covalent radius decrease going across a period?

A

As you from from left to right, you add a proton to the nucleus and an electron to the outermost energy level

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15
Q

What happens to covalent radius going down a group?

A

The covalent radius increases

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16
Q

Why does covalent radius increase going down a group?

A

An extra outer electron shell is added every time making it further away from the nucleus

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17
Q

What happens to atomic size across a period?

A

The atomic size decreases

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18
Q

Why does atomic size decrease going across a period?

A

A proton is added each time, increasing the nuclear charge

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19
Q

What happens to atomic size going down a group?

A

The atomic size increases

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20
Q

Why does atomic size increase going down a group?

A

This is due to an outer shell of electrons being added each time

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21
Q

What is ionisation energy?

A

The energy required to remove one mole of electrons from one mole of atoms in the gaseous state

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22
Q

What happens to ionisation energy going across a period?

A

The ionisation energy increases

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23
Q

Why does ionisation energy increase across a period?

A

Due to increasing nuclear charge

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24
Q

What happens to ionisation energy going down a group?

A

The ionisation energy decreases

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25
Why does ionisation energy decrease going down a group?
Due to the electrons being further away from the nucleus and the shielding effect from inner electrons
26
What happens in the second ionisation energy?
Negative electrons are being removed from positive ions rather than neutral atoms
27
What is electronegativity?
A measure of the attraction an atom has for the electrons involved in a bond
28
What happens to electronegativity going across a period
The electronegativity increases
29
Why does the electronegativity increase going across a period?
Because the charge in the nucleus increases across a period
30
What happens to the electronegativity going down a group?
The electronegativity decreases
31
Why does electronegativity decrease going down a group?
Because the atoms have a bigger radius
32
What is the electronegativity difference for pure covalent?
< 0.4
33
What is the electronegativity difference for polar covalent?
Between 0.4 and 1.8
34
What is the electronegativity difference for ionic?
> 1.8
35
When does polar covalent arise?
When there is a difference in electronegativity between atoms in a bond
36
How is a negative dipole formed?
The atom with the higher electronegativity will pull the bonding electrons closer to itself
37
Where does hydrogen bonding occur?
Between molecules of compounds that are highly electronegative like N, O or F
38
Why do boiling points increase when molecular size increases?
Because of Van der Waal’s forces
39
What does hydrogen bonding in solid water give rise to?
An open structure
40
Why is ice less dense than liquid water and float to the top?
Because it has an open structure
41
When does a permanent dipole exist?
If the molecule has a permanent uneven distribution of charge
42
How is solubility of a substance determined?
By it’s bonding
43
What can polar substances dissolve in?
Other polar substances
44
What can non-polar substances dissolve in?
Other non-polar substances
45
What does diatomic molecule bonding exist in?
Diatomic molecules
46
What are the forces like in diatomic molecules?
Forces within the molecules are strong covalent bonds whereas forces between the molecules are weak Van der Waal’s
47
What are melting and boiling points like in diatomic molecules?
Melting and boiling points are low
48
What are examples of elements that contain diatomic molecules?
- Hydrogen - Nitrogen - Fluorine - Oxygen - Chlorine
49
What holds the atoms together in covalent network bonding?
Strong covalent bonds
50
What are melting and boiling points like in covalent network?
Very high melting and boiling points
51
What are examples of elements that have covalent network?
- Boron - Carbon - Silicon
52
What are examples of elements with covalent molecular?
- Phosphorus - Sulfur
53
What do covalent molecular molecules produce?
Larger Van der Waal’s forces than diatomic molecules because they have many electrons
54
What does the stonger temporary dipoles do to the melting and boiling points in covalent molecular substances?
Raises melting and boiling points making them solids at room temperature
55
What does delocalised electrons do to metallic bonding elements?
Makes them good conductors
56
What are examples of metallic bonding elements?
- Lithium - Sodium - Magnesium - Potassium - Calcium
57
What exists between the atoms in monatomic elements?
Weak Van der Waal’s forces
58
What type of melting and boiling points are in monatomic elements?
Very low melting and boiling points
59
What are examples of monatomic elements?
- Helium - Neon - Argon
60
How are London Dispersion Forces caused?
By the uneven distribution of the constantly moving electrons around the nuclei of the atoms. This then causes the formation of temporary dipoles on the atoms. The atoms then attract each other
61
Why do London Dispersion Forces increase?
Because as you go down the group, the number of electrons increases
62
What is oxidation?
Loss of electrons
63
What is reduction?
Gain of electrons
64
What is an oxidising agent?
A chemical species that brings about an oxidation in something else but is itself reduced
65
What do oxidising elements have and why?
Elements that are good oxidising agents generally have high electronegativity as they need to be good at accepting electrons from other substances
66
What is a reducing agent?
A chemical species that brings about a reduction in something else but is itself oxidised
67
What do reducing elements usually have and why?
Elements that are good reducing agents generally have low ionisation energies as they need to be good at donating electrons to other substances
68
What is chromatography?
A useful technique to separate molecules in a mixture
69
What is chromatography used for?
Fingerprinting, defecting bombs in airports, identify drugs and alcohol
70
How do you work out Rf value?
Rf = distance from the bass line to the spot / distance from the bass line to the solvent front
71
What is the stationary phase?
Chromatography paper
72
What is the mobile phase?
Suitable solvent (Water, Ethanol, Organic Solvent)
73
What is the separation method?
As the solvent moves up the paper, it dissolves the sample. The more soluble the sample, the further up the paper it will travel

Higher Chemistry (3 decks)

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