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AS Chemistry > 2. Bonding > Flashcards

2. Bonding Flashcards

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1
Q

How are ionic compounds formed?

A

Transfer of electrons from metal atoms to non metal atoms

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2
Q

Define an ion

A

Atoms which have lost or gained electrons to gain an overall charge

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3
Q

How do metals form ions

A

Lose electrons to become positively charged

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4
Q

How do non-metals form ions

A

Gain electrons to become negatively charged

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5
Q

Define a cation

A

Positive ions

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6
Q

Define an anion

A

Negative ions

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7
Q

Define an ionic bond

A

The electrostatic attractive between oppositely charged ions

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8
Q

Why do atoms lose/gain electrons?

A

To get a noble gas configuration

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9
Q

Define ionic lattice

A

The structure where ions are held in a 3D framework

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10
Q

Define lattice

A

Regular repeated 3D arrangement of ions, atoms or molecules in a solid

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11
Q

What causes the crystalline nature of ionic compounds?

A

The regular pattern of the ions

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12
Q

Define decrepitation

A

The cracking noise that ionic compounds make when they’re heated caused by the crystalline structure

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13
Q

Why do ionic compounds have high melting/boiling points?

A

Because lots of energy is required to overcome the strong electrostatic forces of attraction between oppositely charged ions

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14
Q

What factors affect the strength of the ionic bond?

A

The size of the ion and the size of the charge

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15
Q

Why are +ve generally smaller?

A

Bc they’ve lost ions so fewer shells and increased effective nuclear charge so shells drawn in closer

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16
Q

Why are -ve ions generally larger?

A

Gain electrons so more repulsion of electrons pushes them further apart
Smaller effective nuclear charge bc more electrons with same no of protons

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17
Q

Are ionic compounds soluble?

A

Most of the time, depending on the strength of the ionic bonds

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18
Q

Why can ionic compounds only conduct electricity in the molten/aqueous state?

A

Because the ions are free to move and carry charge in these states

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19
Q

Define a covalent bond

A

One or more shared pairs of electrons between two non-metal atoms

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20
Q

Define a single covalent bond

A

One shared pair of electrons

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21
Q

Define a double covalent bond

A

Two pairs of shared electrons

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22
Q

Define a triple covalent bond

A

Three pairs of shared electrons

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23
Q

What is a bonding pair of electrons?

A

Pair of electrons that are shared between two atoms

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24
Q

What is a lone pair of electrons?

A

An unshared pair of electrons that are not involved in bonding

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25
Q

Define a coordinate bond

A

A bond which has a shared pair of electrons where both electrons are supplied by one atoms

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26
Q

How are ions in metals arranged?

A

In layers with the outer shell electrons not bound to an atom

27
Q

Define delocalised electrons

A

Electrons that are not bound to one particular atom

28
Q

What is a metallic bond?

A

The electrostatic attraction between delocalised electrons and positive ions in the lattice

29
Q

Why can metals conduct electricity?

A

Delocalised electrons are free to move throughout the structure and can carry charge

30
Q

Why can metals conduct heat?

A

Delocalised electrons can move and carry heat energy throughout the metal

31
Q

Why are metals malleable and ductile/

A

Because the layers can slide over each other without disrupting the bonding

32
Q

Why do metals have high densities?

A

Because the positive ions are packed tightly together

33
Q

Why do metals have high melting points?

A

Hella energy required to overcome the strong forces of attraction between the positive ions and delocalised electrons

34
Q

Why do some metals have a higher melting point than others?

A

Because they have more delocalised electrons and they may have smaller ions so they’re packed more tightly together

35
Q

What are molecular covalent substances?

A

Substances which exist as single molecules

36
Q

What are molecular covalent crystals?

A

Molecular covalent substances which form crystalline structures

37
Q

Why are molecular covalent crystals brittle?

A

Because they dont have strong bonds holding them together

38
Q

Why do molecular covalent crystals not conduct electricity?

A

Bc they don’t have any charged particles to carry charge

39
Q

Define an allotrope

A

Different forms of the same element in the same physical state

40
Q

What are the two allotropes of carbon?

A

Diamond and graphite

41
Q

Why is diamond hard?

A

Many strong covalent bonds which require hella energy to overcome

42
Q

Why does diamond have a high melting point?

A

Strong covalent bonds that require hella energy to overcome

43
Q

In diamond how many bonds does each carbon form?

A

4

44
Q

Why does diamond not conduct electricity?

A

No charged particles which can move and carry charge

45
Q

In graphite how many bonds does each carbon form?

A

Three

46
Q

Why can graphite conduct electricity?

A

Delocalised electrons which can move and carry charge

47
Q

Why does graphite have high melting point

A

Strong covalent bonds which require a lot of energy to break

48
Q

Why is graphite a good lubricant?

A

Layers can slide over each other because of weak forces in between layers

49
Q

What are the types of crystalline structures?

A

Ionic, metallic, macromolecular and molecular

50
Q

Why do lone pairs have a greater repulsive force?

A

Because they’re closer to the central atom

51
Q

Linear: bonding pairs and angles

A

2 bonding pairs with 180 degrees between them

52
Q

Trigonal planar: bonding pairs and angles

A

3 bonding pairs with 120

53
Q

Tetrahedral: bonding pairs and angles

A

4 pairs with 109.5

54
Q

Trigonal bipyramidal: bonding pairs and angles

A

5 pairs with 90 and 120

55
Q

Octahedral: bonding pairs and angles

A

6 bonding pairs all 90

56
Q

Pyramidal: bonding pairs and angles

A

3 bonding pairs, one lone pair - 107

57
Q

Bent: bonding pairs and angles

A

2 bonding pairs, 2 lone pairs - 104.5

58
Q

Define electronegativity

A

Power of an atom to attract the pair of electrons in a covalent bond

59
Q

What factors affect electronegativity?

A 
Atomic radius (distance between pair of electrons and nucleus)
Nuclear charge
Shielding
60
Q

Why does electronegativity increase across a period?

A

Decrease in atomic radius so stronger attraction between nucleus and bonding pair
Increase in nuclear charge so more attraction

61
Q

Why does electronegativity decrease down a group?

A

Increased atomic radius and weaker nuclear charge b more shielding so weaker attraction between nucleus and bonding pairs

62
Q

How are induced dipoles formed?

A

Electrons are in constant motion so they may be distributed more on one side of the molecule creating a temp dipole

63
Q

What is the relationship between van der Waals forces and the size of the molecule?

A

Larger molecules have more electrons so greater induced dipoles so greater van der Waals forces between molecules

64
Q

How are hydrogen bonds formed?

A

Occur between delta positive charge of a H in a bond and the lone pair of electrons of an O, N or F of another molecule or atom

AS Chemistry (14 decks)

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