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chemistry > section 3 ; bonding > Flashcards

section 3 ; bonding Flashcards

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1
Q

what is an ion

A

formed when electrons are transferred from one atom to another. Simplest ions are single atoms that have lost or gained 1,2,3 electrons to have a full outer shell

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2
Q

what is ionic bonding

A

electrostatic attraction holds positive and negative ions together (very strong). When atoms are held together in a lattice it is called ionic bonding. When oppositely charged ions come together and form ionic bonds you get an ionic compound

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3
Q

what substances ionically bond

A

metal with non metal

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4
Q

what is the structure of ionic bonding

A

giant lattice structure of pos and neg ions - ions held together by ionic bonds (ES attraction between pos and neg ions)

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5
Q

what are the melting points and boiling points of ionic bonds

A

HIGH strong electrostatic forces of attraction (ionic bonds) need to be broken higher b and m points happen depending on size and charge of ions

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6
Q

are ionic compounds conductive

A

YES when molten or dissolved ions are free to move and carry charge

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7
Q

are ionic compounds soluble

A

tend to dissolve in water water molecules are polar - part of the molecule has a small negative charge and the other bits have a small positive charge, water molecules pull the ions away from lattice causing it to dissolve

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8
Q

what is a giant lattice

A

a regular structure = lattice basic same unit (lattice) repeated over and over again = giant lattice

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9
Q

how does covalent bonding occur

A

two atoms share electrons so both got full outer shells of electrons. A single covalent bond contains a shared pair of electrons, both the pos nuclei are attracted electrostatically to the shared electrons

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10
Q

what is a simple molecular (simple covalent) compound

A

elements are diatomic atoms eg. N2 O2 compounds are non metals with non metals

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11
Q

what bonding is involved with simple molecular

A

covalent

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12
Q

are B and M points low or high with simple molecular

A

generally low - weak intermolecular forces between molecules , easy to break (NOT breaking strong covalent bonds between atoms) M and B increase depending on size of molecule/ Mr and the type of intermolecular

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13
Q

are simple molecular compounds good conductors

A

generally no - no delocalised e to carry the charge

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14
Q

are simple molecular compounds soluble

A

no

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15
Q

what is a giant covalent structure

A

lattice structure in which all atoms are joined to others by covalent bonds eg. graphite graphene diamond silicon dioxide - type of crystal structure - sometimes called macromolecular structures

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16
Q

what bonding holds together giant covalent structures

A

covalent

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17
Q

what are the B and M points of giant covalent structures like

A

VERY HIGH - breaking strong covalent bonds between atoms depends on number and strength of covalent bonds

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18
Q

are giant covalent good conductors

A

no EXECPT graphite and graphene

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19
Q

what is a metallic giant lattice

A

lattice structure of metal ions with outer shell electrons free to move through the structure. Strong electrostatic force of attraction between positive ions and negative delocalised electrons. closely packed positive ions in a sea of delocalised electrons (metallic bonding) - between metals

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20
Q

what are the M and B points like for metallic bonding

A

HIGH strong electrostatic force of attraction between positive ions and sea of delocalised electrons increases/depends on n of delocalised electrons

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21
Q

can giant metallic lattices conducted electricity

A

YES because of sea of delocalised electrons able to carry the charge

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22
Q

are giant metallic lattices soluble

A

no

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23
Q

what is the strength of ionic bonds like

A

the smaller the ions and the greater the charge on the ions the stronger the attraction between pos and neg ions (usually)

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24
Q

what is the strength of covalent bonds like

A

the shorter the bond the stronger the bond (usually) triple bonds are stronger than double which are stronger than single

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25
Q

what is the strength of metallic bonds like

A

the smaller the metal ions, the greater charge of the ions and the more delocalised outer shell electrons there are the stronger the attraction between ions and electrons (usually) and therefore higher melting and boiling points

26
Q

explain charge density in reference to ionic bonding

A

how spread out the charge is over large area (atom size) the charge is more ‘dilute’ eg. if two different atoms have same charge but one is larger, the smaller one has stronger bonds and therefore higher melting and boiling points.

27
Q

explain how single, double and triple bonds are formed with reference to covalent bonds

A
  • two shared e-s = single bond - four shared e-s = double bond - six shared e-s = triple bond
28
Q

what is hybridisation

A

when atoms merge energy levels to allow room for more electrons

29
Q

explain how to draw covalent molecule diagrams

A

show free electrons as ‘X’

remember to highlight double or single bonds

30
Q

what are co-ordinate (dative covalent) bonds

A

where BOTH the electrons come from the same species

  • often drawn with a “->” rather than a “-“ showing the direction in which the electrons are donated
  • however once formed they are identical to covalent bonds
31
Q

how to draw allotropes of carbon in exam

A
32
Q

why do graphite and graphene conduct electricity

A

4th delocalised e- in between layers , free to move and pass current

33
Q

define toughness in terms of chemistry

A

ability of a material to absorb energy without fracturing

34
Q

why is graphite a soft metal

A

layers can slide

  • weak vander waals forced between layers
35
Q

why is ice less dense than water

A

ice is in a regular lattice rather than a random arragement like water

  • in ice there are hydrogen intermolecular forces which hold the molecules in the lattice far apart (more than in water)
  • this means ice is less dense
36
Q

why would an ionic compound be soluble

A

depending on how polarised the molecules are (water attracks them)

37
Q

what are van der waals forces

A

an intermolecular force present in all molecular substances

  • also known as induced dipole-dipole forces or london dispersion forces
38
Q

why do van der waals forces exist

A

because electrons are constantly moving around and uneven electron distribution at any given time

  • causes a temporary dipole within an molecule which induces a temporary dipole in an neighbouring molecule
  • attraction between these molecules , temporary induced dipole-dipole attraction
  • the more electrons there are, the larger the dipole
39
Q

how does van der waals forces increase

A

the bigger the molecule

-straight chain have more than branched due to increased surface contact area for bonds to form

40
Q

what is a permanent dipole-dipole attraction

A

permanent dipole-dipole attraction between polar molecules

41
Q

what is a hydrogen bond

A
  • special case of permanent dipole-dipole attraction where a hydrogen atom is attracted to a very electronegative atom (eg. F, O, N)
  • simply a strong intermolecular force NOT a bond
42
Q

show and explain how to draw hydrogen bonds between molecules

A
  • always show all lone pairs of electrons , all partial positive and negative charges and dotted line representing hydrogen bond
43
Q

describe the different strengths of intermolecular forces

A

H- bonding > permanent DD > wav der waals

44
Q

what determines the shape of a molecule

A
  • bonding pairs (the two shared electrons in a covalent bond)
  • lone pairs (two electrons in a pair not involved in bonding)

These electron pairs will repel each other as far as possible

The shape of any molecule or ion is a consequence of the number of electron pairs which repel each other as far as possible

45
Q

explain the five basic shapes

A
46
Q

explain how to work out the shape for molecules and ions (single bonds only)

A

1) work out how many outer shell electrons are on the central atom (accounting for any charge on the species)
2) work out how many atoms are bonded to the central atom (= the number of bonding pairs)
3) work out how many lone pairs there (half difference between 1 and 2)

47
Q

how do lone pairs of electrons affect the angles between bonded atoms to the main atom

A

reduces the angle by 2.5 degrees

48
Q

show the other shapes you must remember for 6 electron pairs

A
49
Q

state the meaning of electronegativity

A

the power of an atom to attract the two electrons in an covalent bond

50
Q

name the factors that affect electronegativity

A

1) proton number (nuclear charge) - the more protons the stronger the attraction between the protons and es - greater electro neg
2) sheilding - the more shells the harder it is for the protons to attract the electrons , lower EN
3) size - smaller the atom the stronger the ESFOA

51
Q

explain the trend of electronegativity down a group

A

decreases as more shells so more sheilding

  • larger atoms too, less attraction
52
Q

explain the trend in electronegativity across a period

A

increases

  • atoms are smaller (smaller atomic radius) and higher number of protons
53
Q

explain what non-polar covalent bonds are

A

when two atoms in a covalent bond have the same electronegativity

  • the two electrons are shared equally
    eg. Cl-Cl
54
Q

what are polar covalent bonds

A

when two atoms in a covalent bond have a different electronegativity

  • the more electronegative atom has a greater share of the two electrons while the less electronegative has a lower share
    eg. H-Cl
55
Q

are C-H bonds counted as being polar in organic molecules

A

no

56
Q

explain bond dipole moment

A

measure of strength and direction of the polarity in the bond , the bigger the difference in EN the bigger the bond dipole moment

57
Q

show shapes for 5 electron pair molecules

A
58
Q

show the shapes for 4 electron pair molecules

A
59
Q

show the shapes of 3 electron pair molecules

A
60
Q

show the shapes of 2 electron pair molecules

A

chemistry (9 decks)

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