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Bonding Flashcards

(53 cards)

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1
Q

How are ions formed?

A

When one or more electrons are transferred from one atom to another

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2
Q

What is ionic bonding?

A

Ions are held together by electrostatic attraction

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3
Q

What are compound ions?

A

Lots of ions that make up groups of atoms with an overall charge

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4
Q

What is a giant ionic lattice?

A

Same basic unit repeated over and over again. Ionic crystals are giant lattices of ions

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5
Q

How does ionic structure explain the behaviour of ionic compounds?

A
  1. Ionic compounds conduct electricity when they’re molten or dissolved - but not when they’re solid
  2. Ionic compounds have high melting points
  3. Ionic compounds tend to dissolve in water
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6
Q

Why do ionic compounds conduct electricity when they’re molten or dissolved - but not when they’re solid?

A

The ions in a liquid are free to move (and they carry a charge).
In a solid the ions are fixed in position by strong ionic bonds

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7
Q

Why do ionic compounds have high melting points?

A

Giant ionic lattices are held together by strong electrostatic forces
It takes loads of energy to overcome these forces, so melting points are very high

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8
Q

Why do ionic compounds tend to dissolve in water?

A

Water molecules are polar - part of the molecule has a small negative charge and other bits have small positive charges
These charged parts pull ions away from the lattice, causing it to dissolve

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9
Q

What are molecules?

A

Groups of atoms bonded together

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10
Q

How are molecules held together?

A

Strong covalent bonds

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11
Q

What does a single covalent bond contain?

A

A shared pair of electrons

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12
Q

What happens in covalent bonding?

A

Two atoms share electrons so they’ve both got full outer shells of electrons
Both the positive nuclei are attracted electrostatically to the shared electrons

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13
Q

What are giant covalent structures?

A

Have a huge network of covalently bonded atoms

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14
Q
A
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15
Q

How many bonds can carbon form?

A

4 strong, covalent bonds

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16
Q

How does the structure of graphite determine its properties?

A
  1. The weak bonds between the layers in graphite are easily broken, so the sheets can slide over each other - graphite feels slippery and is used as a dry lubricant and in pencils
  2. The ‘delocalised’ electrons in graphite aren’t attached to any particular carbon atom and are free to move along the sheets carrying a charge. So graphite is an electrical conductor
  3. The layers are quite far apart compared to the length of the covalent bonds, so graphite has a very high melting point - sublimes at over 3900K
  4. Graphite is insoluble in any solvent. The covalent bonds in the sheets are too strong to break
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17
Q

What shape is diamond?

A

Tetrahedral

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18
Q

How do the strong covalent bonds of diamond determine its properties?

A
  1. High MP - sublimes at over 3900K
  2. Extremely hard - used in diamond-tipped drills and saws
  3. Vibrations travel easily through the stiff lattice, so goof thermal conductor
  4. Can’t conduct electricity - all the outer electrons are held in localised bonds
  5. Like graphite, diamond won’t dissolve in any solvent
  6. You can ‘cut’ diamond to form gemstones. Its structure makes it refract light a lot, which is why it sparkles
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19
Q

What is dative bonding?

A

Where both electrons come from one atom

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20
Q

How do bonding pairs and lone pairs exist?

A

As charge clouds

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21
Q

What is a charge cloud?

A

An area where you have a really big chance of finding an electron pair
Electron charge clouds repel each other

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22
Q

What does molecular shape depend on?

Study These Flashcards
A

Electron pairs around the central atom

23
Q

How do you use number of electron pairs to predict the shape of a molecule?

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A
  1. Work out which one is central atom
  2. Use periodic table to work put number of electrons in outer shell of central atom
  3. Add one to this number for every atom that the central atom is bonded to. If ion, add one for each negative charge / subtract one for each positive charge
  4. Divide by two to find number of electron pairs on the central atom
  5. Compare number of electron pairs to number of bonds to find the number of lone pairs and the number of bonding pairs on the central atom
24
Q

How many bonding pairs, lone pairs and the bond angle for linear?

Study These Flashcards
A

BP: 2
LP: 0
Angle: 180

25
How many bonding pairs, lone pairs and the bond angle for trigonal planar?
BP: 3 LP: 0 Angle: 120
26
How many bonding pairs, lone pairs and the bond angle for tetrahedral?
BP: 4 LP: 0 Angle: 109.5
27
How many bonding pairs, lone pairs and the bond angle for trigonal pyramidal?
BP: 3 LP: 1 Angle: 107
28
How many bonding pairs, lone pairs and the bond angle for bent?
BP: 4 LP: 2 Angle: 104.5
29
How many bonding pairs, lone pairs and the bond angle for trigonal bipyramidal?
BP: 5 LP: 0 Angle: 90, 120
30
How many bonding pairs, lone pairs and the bond angle for seesaw?
BP: 4 LP: 1 Angle: 87, 102
31
How many bonding pairs, lone pairs and the bond angle for T-shaped?
BP: 6 LP: 2 Angle: 88
32
How many bonding pairs, lone pairs and the bond angle for octahedral?
BP: 6 LP: 0 Angle: 90, 90
33
How many bonding pairs, lone pairs and the bond angle for square planar?
BP: 4 LP: 2 Angle: 90
34
What is electronegativity?
An atom's ability to attract the electron pair in a covalent bond is called electronegativity
35
What makes a bond polar?
In a covalent bond between two atoms of different electronegativities, the bonding electrons will be pulled towards the more electronegative atom
36
What makes a bond non-polar?
The atoms have equal electronegativities, so the electrons are equally attracted to both nuclei
37
What causes a permanent dipole?
In a polar bond, the difference in electronegativity between the two atoms
38
What is a dipole?
A difference in charge between the two atoms caused by a shift in electron density in the bond
39
What increases polarity?
The greater the difference in electronegativity between the atoms
40
What are the three intermolecular forces?
Induced dipole-dipole / Van der Waals forces Permanent dipole-dipole forces Hydrogen bonding
41
Where are Van der Waals forces found?
Between all atoms and molecules
42
What makes Van der Waals forces stronger?
Larger molecules have larger electron clouds meaning stronger VdW forces
43
How does shape of molecules affect strength of VdW forces?
Long, straight molecules can lie closer together than branched ones - the closer together two molecules get, the stronger the forces between them ar
44
When does hydrogen bonding occur?
When hydrogen is covalently bonded to fluorine, nitrogen or oxygen
45
Why are hydrogen bonds so strong?
F, N and O are very electronegative So they draw bonding electrons away from the hydrogen atom The bond is so polarised, and hydrogen has such a high charge density that the hydrogen atoms form weak bonds with lone pairs of electrons on the fluorine, nitrogen or oxygen atoms of other molecules
46
How does hydrogen bonding effect the properties of substances?
Substances with hydrogen bonds have higher boiling and melting points than other similar molecules because of the extra energy needed to break the hydrogen bonds As liquid water cools to form ice, the molecules make more hydrogen bonds and arrange themselves into a regular lattice structure In this regular structure, the H2O molecules are further apart on average than the molecules in liquid water - so ice is less dense than liquid water
47
Why do metals have giant metallic lattice structures?
1. The outermost shell of electrons of a metal atom is delocalised - the electrons are free to move about the metal. This leaves a positive metal ion 2. The positive metal ions are attracted to the delocalised negative electrons. They form a lattice of closely packed positive ions in a sea of delocalised electrons
48
What is metallic bonding?
Positive metal ions attracted to delocalised negative electrons forming a lattice of closely packed positive ions in a sea of delocalised electrons
49
How does metallic bonding explain the properties of metals?
1. Metals have high melting points because of the strong electrostatic attraction between the positive metal ions and the delocalised sea of electrons 2. The number of delocalised electrons per atom affects the melting point. The more there are, the stronger the bonding will be and the higher the melting point 3. The delocalised electrons can pass kinetic energy to each other, making metals good thermal conductors 4. Metals are good electrical conductors because the delocalised electrons can move and carry a charge 5. Metals are insoluble (except in liquid metals), because of the strength of the metallic bonds
50
Why do simple covalent compounds have relatively low melting and boiling points?
To melt or boil a simple covalent compound you only have to overcome the intermolecular forces that hold the molecules together You don't need to break the much stronger covalent bonds that hold the atoms together in the molecules
51
When will a solid substance conduct electricity?
If it contains charged particles that are free to move
52
What are melting and boiling points determined by in a solid?
By strength of attraction between its particles
53
What determines how soluble a solid substance is in water?
Type of particles it contains Water is a polar solvent, so substances that are polar or charged will dissolve in it well, whereas non-polar or uncharged substances won't

Chemistry Physical (5 decks)

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