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4... Flashcards

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1
Q

ionic bonds

A
  • transfer of electrons from a metallic element to a non-metallic element
  • cation and anions are attracted to each other
  • electrostatic attraction - very strong
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2
Q

cations

A

metals loose electrons from their valance shell forming positively charged cations

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3
Q

anions

A

non-metal atoms gain electrons forming negatively charged anions

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4
Q

lattice structure

A
  • evenly distributes crystalline structure
  • ## regular repeating patter so the + charges cancel out - charges
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5
Q

melting and boiling points of ionic compounds

A
  • ionic compounds are strong
  • high melting and boiling points
  • strong electrostatic forces between ion
  • melting and boiling point increases with charge density as they have higher electrostatic forces
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6
Q

solubility of ionic compounds

A
  • soluble in water
  • they form ion - dipole bonds
  • can conduct electricity when molten or in a solution
  • ions can move freely around
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7
Q

covalent bonds

A
  • occurs between two non-metals
  • electrons are shared
  • each atom provide one of the electrons in the bond
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8
Q

covalent bonding electronegativity difference

A

<1.0

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9
Q

polar covalent electronegativity difference

A

1.0-2.0

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10
Q

ionic electronegativity difference

A

> 2.0

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11
Q

coordinate bonds

A
  • both electron in the bond are from the same atom
  • some molecules have a lone pair of electron that can be donates to form a bond with an electron-deficient atom
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12
Q

bond length

A
  • internuclear distance of two covalnetly bonded atoms
  • as the strength of the covalent bond increases, the length of bond decreases

triple bond<double bond<single bond

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13
Q

non polar bonds

A

two atoms have the same electronegativity

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14
Q

polar bonds

A

different electronegativites, electron drawn towards the more electronegative atom
less electronegative - partial charge of δ+
more electronegative - partial charge of δ-

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15
Q

dipole moment

A
  • measure of how polar a bond is
  • direction of the dipole moment is shown by the following sign in which the arrow points to the partially negatively charged end of dipole
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16
Q

steps for drawing a lewis structures

A
  1. count the total number of valance
  2. draw the skeletal structure to show how many atoms are linked to each other
  3. use pair of crosses or dot/cross to put an electron pain in each bon between atome
  4. add more electron pairs to complete the octets around the atoms
  5. if there are not enough electrons to complete the octets, form double/triple bonds
  6. check the total number of electron in the finished structure is equual to the total number of valance electrons
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17
Q

exceptions to the octet rule

A

H, Li, Be, B, and Al

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18
Q

delocalized electrons

A

electrons in a molecule, ion and solid metal that are not associated with a single atom or one covalent bond
e.g. double bond capable of migrating

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19
Q

electron and molecular geometry when there is 2 bonded electron pairs, and 0 lone pair

A

electron domain geometry: linear
molecular geometry: linear
bond angle: 180

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20
Q

electron and molecular geometry when there is 3 bonded electron pairs, and 0 lone pair

A

electron domain geometry: trigonal planar
molecular geometry: trigonal planar
bond angle: 120

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21
Q

electron and molecular geometry when there is 2 bonded electron pairs, and 1 lone pair

A

electron domain geometry: trigonal planar
molecular geometry: bent linear
bond angle: 119

22
Q

electron and molecular geometry when there is 4 bonded electron pairs, and 0 lone pair

A

electron domain geometry: tetrahedral
molecular geometry: tetrahedral
bond angle: 109.5

23
Q

electron and molecular geometry when there is 3 bonded electron pairs, and 1 lone pair

A

electron domain geometry: tetrahedral
molecular geometry: trigonal pyramid
bond angle: 107

24
Q

electron and molecular geometry when there is 2 bonded electron pairs, and 2 lone pair

A

electron domain geometry: tetrahedral
molecular geometry: bent linear
bond angle: 104.5

25
Q

molecular polarity

A

molecule is able to have polar bonds but not be overall polar
molecules that are symmetrical are unlikely to be polar

26
Q

characteristics of a diamond

A

melting/boiling point: very high
electrical conductivity: non-conductor
appereance: transparent crystals
special characteristic: hardest known naturally occuring substance

27
Q

characteristics of a graphite

A

melting/boiling point: very high
electrical conductivity: good
appereance: grey - black solid
special characteristic: soft and slippery

28
Q

characteristics of a Buckminster-fulleren (C60)

A

melting/boiling point: low
electrical conductivity: semi-conductor
appereance: yellow solid
special characteristic: light and soft

29
Q

characteristics of a graphene

A

melting/boiling point: very high
electrical conductivity: very good
appereance: transperent sheets
special characteristic: very strong and flexible

30
Q

characteristics of a silicon dioxide

A

melting/boiling point: very high
electrical conductivity: non-conductor
appereance: transparent sheets
special characteristic: produces electric charge from mechanical stress

31
Q

Diamond

A
  • Diamond is a giant lattice of carbon atoms
  • Each carbon is covalently bonded to four others in a tetrahedral arrangement with a bond angle of 109.5
  • The result is a giant lattice with strong bonds in all directions
32
Q

Graphite

A
  • In graphite, each carbon atom is bonded to three others in a layered structure
  • The layers are made of hexagons with a bond angle of 120o
  • The spare electron is delocalised and occupies the space in between the layers
  • All atoms in the same layer are held together by strong covalent bonds, and the different layers are held together by weak intermolecular forces
33
Q

Buckminsterfulleren (C60)

A

Buckminsterfullerene is one type of fullerene, named after Buckminster Fuller, the American architect who designed domes like the Epcot Centre in Florida
It contains 60 carbon atoms, each of which is bonded to three others by single covalent bonds
The fourth electron is delocalised so the electrons can migrate throughout the structure making the buckyball a semi-conductor

34
Q

Graphene

A

Some substances contain an infinite lattice of covalently bonded atoms in two dimensions only to form layers. Graphene is an example
Graphene is made of a single layer of carbon atoms that are bonded together in a repeating pattern of hexagons
Graphene is one million times thinner than paper; so thin that it is actually considered two dimensional

35
Q

Silicon(IV)oxide

A

Silicon(IV)oxide is also known as silicon dioxide, but you will be more familiar with it as the white stuff on beaches!
Silicon(IV)oxide adopts the same structure as diamond - a giant structure made of tetrahedral units all bonded by strong covalent bonds
Each silicon is shared by four oxygens and each oxygen is shared by two silicons
This gives an empirical formula of SiO2

36
Q

London (dispersion) forces

A
  • electrons in atoms are not static
  • a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles.
  • This force is sometimes called an induced dipole-induced dipole attraction.
  • ## present between all atoms and molecules
37
Q

what does the strenght of london dispersion forces depend on?

A

number of electrons:
- greater number of electrons in a molecule = greater frequency and magnitude of temporary dipole
- melting and boiling points are larger
surface area of molecules:
- more contact it will have with adjacent molecules
- greater london dispersion forces
- higher melting anf boiling point

38
Q

dipole - dipole attractions

A

permanent dipole-dipole bonding
- slight increase in strength of intermolecular attractions
- increased boiling point
-

39
Q

hydrogen bonding

A

strongest type of intermolecular forces
- species with O or N or F atom with an available lone pair of electrons
- a hydrogen attached to the O or N or F
- the H becomes so δ+, that it can form a bond with the lone pair of an N or O atom

40
Q

What influences the physical properties of covalent compounds

A
  • physical properties of molecular covalent compounds are largely influenced by their intermolecular forces
  • if you know the type of intermolecular forces present you can predict the physical properties like melting and boiling point, solubility and conductivity.
41
Q

melting and boiling point in relation to intermolecular forces

A

stronger the forces = higher melting and boiling point

42
Q

what affects the strength of intermolecular forces?

A
  • size of molecule
  • ## increase in the polarity of molecule
43
Q

solubility of polar compounds

A

polar dissolves in polar solvent as a result of dipole - dipole interactions or the formation fo hydrogen bonds between solute and solvent

44
Q

solubility of non polar compounds

A

dissolve in non-polar solvents, they form dispersion forces between the solvents and the solute

45
Q

how are metals atoms packed together?

A

lattice structure

46
Q

what does the lattice structure of allow for the electrond to do

A
  • the electrons in their outer shells are free to move throughout the structure
  • free-moving electrons are ‘delocalised’
  • when electrons are delocalised, the metal atoms become positively charged
  • positive charges repel each other and keep the neatly arranged lattice in place
  • ## positive metal centers are suspended in a ‘sea’ of delocalised electrons
47
Q

properties of metals

A
  • malleable
  • metal layers can slide
  • strong and hard
  • high melting and boiling points
48
Q

conductivity of metals

A
  • ## can conduct electricty when in the solid or liquid state
49
Q

what affects the strength of metallic bonds

A

charge on the metal ion
- greater the charge = greater number of delocalised elecctrons, greater charge difference
- stronger electrostatic attractions, stronger metallic bond

radius of the metal ion
- smaller ionic radius means stronger metallic bonds

50
Q

melting points of metals across a period

A

the stronger the metallic bonding, the more nergy is need to break the metallic lattice and so the higher melting point
- as you go across the higher melting point

51
Q

melting points down the group

A

as you go down the group, the attraction between valance electrons and the metalic lattice decreases
decrease in metling point

52
Q

what are alloys

A
  • Alloys are mixtures of metals, where the metals are mixed together physically but are not chemically combined
  • They can also be made from metals mixed with nonmetals such as carbon
  • Ions of the different metals are spread throughout the lattice and are bound together by the delocalized electrons
  • It is possible to form alloys because of the non-directional nature of the metallic bonds

Chem (13 decks)

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