A-LEVEL CHEM > BONDING > Flashcards

BONDING Flashcards

(34 cards)

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1
Q

State how ions form and why they attract each other

A
  • Ions form when atoms lose or gain electrons
  • Since electrons are negatively charged, an atom that loses one or more electrons will become positively charged
  • An atom that gains one or more electrons will become negatively charged
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2
Q

What’s ionic bonding

A

Bonding that involves electrostatic attraction between oppositely charged ions in a lattice

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3
Q

Formula of sulfate ion

A

SO₄²-

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4
Q

Formula of hydroxide ion

A

OH¯

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5
Q

Formula of nitrate ion

A

NO3-

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6
Q

Formula of carbonate ion

A

HCO3-

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7
Q

Formula of ammonium ion

A

NH₄⁺

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8
Q

State the properties of ionically bonded compounds

A
  • High melting and boiling points > ions in giant ionic lattice are held by strong electrostatic forces of attraction acing in all directions> difficult to overcome
  • Conduct electricity when molten or dissolved in solution > ions of the compound are able to move and carry charge. Unable to conduct when solid as ions are fixed in place so unable to carry charge
  • Soluble in water > both ionic compounds and water molecules are partially charged molecules (polar). > means the partial charges of water break apart the ionic lattice, pulling oppositely charged ions apart and the ionic compound to dissolve.

-Always solids at room temperature > have giant structures and therefore high melting temperatures > in order to melt an ionic compound, a lot of energy is required to break up the lattice of ions

  • Brittle and shatter easily when given a sharp blow > they form a lattice of alternating positive and negative ions
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9
Q

Describe the structure of ionically bonded compounds

A

The structure of this ionic compound is a giant lattice with oppositely charged ions and strong electrostatic forces (attractive forces between ions) between them

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10
Q

Describe a covalent bond

A
  • Forms between 2 non-metals
  • Atoms share some of their outer electrons so that each atom has a stable noble gas arrangement
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11
Q

Describe a co-ordinate bond (dative covalent)

A
  • Bonds contains a shared pair of electrons with both electrons supplied by one atom
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12
Q

Describe the properties of covalently bonded molecules

A
  • Substances composed of molecules are gases, liquids, or solids with low melting points > the strong covalent bonds are only between the atoms within the molecules > there’s only a weak attraction between the molecules so the molecules don’t need much energy to remove the bonds
  • Poor electrical conductors > molecules are neutral overall > no charged particles to carry current
  • When dissolved in water, they remain as molecules > solutions don’t conduct electricity > no charged particles
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13
Q

Define electronegativity

A

measure of tendency of an atom to attract a bonding pair of electrons in a covalent bond

-greater the e.n. of an atom the more it attracts electrons towards it

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14
Q

Factors that affect electronegativity of an atom

A

-atomic charge
-distance from nucleus (atomic nucleus)
-electron sheilding

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15
Q

Non-polar covalent bonds

A

no difference between electronegativities e.g. Cl-Cl

-symmetrical molecules are non-polar even though they contain polar bonds

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16
Q

Polar covalent bonds

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A

Small difference between electronegativities e.g. H-Cl

17
Q

Ionic bond

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A

Large difference in electronegativities

18
Q

Difference between intermolecular and intramolecular forces

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A

Intermolecular - attractive forces between neighbouring molecules

intramolecular -bonds that act WITHIN one molecule

19
Q

Linear shape

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A
  • Total number of electron pairs = not applicable
  • 1 bonding pair of electrons
  • lone pairs not applicable
  • bond angle = 180
    -e.g. H2 and HCL
20
Q

Linear Shape (2)

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A
  • Total number of electron pairs = 2
  • 2 bonding pair of electrons
  • 0 lone pairs
    -180
    -e.g. C02 HCN
21
Q

Trigonal Planar

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A
  • 3 electron pairs
  • 3 bonding pairs
  • 0 lone pairs
    -120
    e.g. Al, Cl3, Bcl3
22
Q

V-shaped/ bend/ non linear

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A
  • 4 electron pairs
    -2 bonding pairs
  • 2 lone pairs
    -104.5
    -e.g. H2O
23
Q

Tetrahedral

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A

-4 electron pairs
-4 bonding pairs
-0 lone pairs
-109.5
-e.g. CH4

24
Q

Triangular pyramid

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A

-4 electron pairs
-3 bonding pairs
-1 lone pair
-107
-e.g. NH3

25
Triangular bi pyramidal
-5 electron pairs -5 bonding pairs -0 lone pairs -90 and 120 -e.g. PCl5
26
Octahedral
-6 electron pairs -6 bonding pairs -0 lone pairs -90 -e.g. SF6
27
How do van der waal forces occur
-uneven distribution of electrons creates a temporary dipole -this induces a dipole in the neighbouring atom/molecule -causes an attraction between diples
28
Why are van der waal forces stronger in larger molecules
more electrons
29
What's permanent dipole dipole
weak attractive force between permanent dipoles in neighbouring polar molecules
30
What's hydrogen bonding
Strong forces of attraction between H nucleus and lone pair of electrons on O,N / F
31
Conditions needed for H bonding to take place
-Bonded to strongly electronegative elements (F,O,N) -electronegative elements MUST have at least a lone pair of electrons
32
Explain hydrogen bonding in hydrogen fluoride
-In HF, the fluorine atom has 3 lone pairs of electrons -F is the most electronegative elements so strongly attracts the lone pair of electrons in a covalent bond -therefore molecule is polar with H atom having a + charge
33
Why's ice less dense than water
-In water, water molecules are constantly moving randomly (molecules far or sometimes close) and so H bonds are constantly being broken and reformed. -As water cools, molecules move more slowly -when reach freezing point, H2O molecules arrange themselves in fixed positions (ice) and this is stabilised by the network of H bonds -H2O molecules in ice are further apart than in liquid as they're in fixed positions
34
Why's ice less dense than water
-In water, water molecules are constantly moving randomly (molecules far or sometimes close) and so H bonds are constantly being broken and reformed. -As water cools, molecules move more slowly -when reach freezing point, H2O molecules arrange themselves in fixed positions (ice) and this is stabilised by the network of H bonds -H2O molecules in ice are further apart than in liquid as they're in fixed positions
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