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chemistry > Topic 2 - bonding and structure CGP > Flashcards

Topic 2 - bonding and structure CGP Flashcards

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1
Q

Define ionic bonding

A

An ionic bond is the strong electrostatic attraction between two oppositely charged ions

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2
Q

What two things can affect te strength of an ionic bond?

A

Ionic charges and ionic radii

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3
Q

How do ionic charges affect the strength of an ionic bond?

A

In general, the greater the charge on an ion, the stronger the ionic bond and therefore, the higher the melting/boiling point

eg. NaF will have a much lower melting point (933C) than CaO (2572C)

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4
Q

How do ionic radii affect the strength of an ionic bond?

A

Smaller ions can pack closer together than larger ions. Electrostatic attraction gets weaker with distance, so small, closely packed ions have stronger ionic bonding than larger ions, which sit further apart. Therefore, ionic compounds with small, closely packed ions have higher melting and boiling points than ionic compounds made of larger ions.

Eg. the ionic radius of Cs+ is greater than that of Na+. NaF has a melting point of 933C, whereas CsF has a melting point of 683C since the Na+ and F- ions can pack closer together in NaF than the Cs+ and F- ions in CsF.

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5
Q

What happens to ionic radius as you go down the group?

A

Ionic radius increases down a group

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6
Q

What happens to the ionic radius in isoelectronic ions?

A

The ionic radius of a set of isoelectronic ions decreases as the atomic number increases.

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7
Q

what structure do ionic compounds form?

A

Ionic compounds form giant ionic lattices

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8
Q

Why do ionic compounds have a high melting point (explaining ionic compound model)

A

As ions are held together by a strong attraction - positive and negative ions are strongly attracted

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9
Q

Why are ionic compounds often soluble in water but not in non polar solvents (explaining ionic compound model)

A

the particles are charged. The ions are pulled apart by polar molecules like water, but not by non polar molecules.

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10
Q

Why do ionic compounds conduct electricity in the molten state but not in the solid state (explaining ionic compound model)

A

Ions in a solid are in a fixed position with strong ionic bonds, therefore no free electrons can carry charge. However, in the liquid/molten state, electrons are free to move (and carry charge)

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11
Q

Why can ionic compounds not be shaped? (explaining ionic compound model)

A

eg. if you try to pull layers of NaCl over each other, you’d get negative chlorine ions directly above other negative chlorine ions (and positive sodium ions directly over each other). The repulsion between these ions would be very strong, so ionic compounds are brittle (they break when they’re stretched or hammered). This supports the lattice model.

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12
Q

explain a proof for ‘the migration of ions is evidence for the presence of charged particles’

A
  • when you electrolyse a green solution of copper (II) chromate (VI) on a piece of wet filter paper, the filter paper turns blue at the cathode (the negative electrode) and yellow at the anode (the positive electrode)
  • copper (II) ions are blue in solution and chromate (VI) ions are yellow. Copper (II) chromate (VI) solution is green because it contains both ions.
  • When you pass a current through the solution, the positive ions move to the cathode and negative ions move to the anode.
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13
Q

What is ionic bonding

A

Ionic bonding is the strong electrostatic attraction between the two positive nuclei and the shared electrons in the bond

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14
Q

In covalent molecules what are the positive nuclei attracted to

A

the area of electron density between the two nuclei (where the shared electrons are)

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15
Q

Explain the repulsion in the covalent molecule and the name of the distance at which these forces are equal (balance out)

A

The two positively charged nuclei repel each other, as do electrons. To maintain the covalent bond these has to be a balance between these forces. The distance between the two nuclei is the distance where the attractive and repulsive forces balance each other. the distance is the BOND LENGTH.

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16
Q

What happens to the bond length if the electron density is higher

A

The higher the electron density between the nuclei (the more electrons in a bond), the stronger the attraction between the atoms, the higher the bond enthalpy and the shorter the bond length.

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17
Q

Explain why a C-C double bond has a shorter bond length than a C-C bond

A

A C–C bond has a greater bond enthalpy and is shorter than a C-C bond. Four electrons are shared in C–C and only two in C-C, so the electron density between the two carbon atoms is greater and the bond is shorter.

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18
Q

What is a dative covalent bond

A

A dative covalent bond is a bond in which one atom donates both electrons to a bond (eg CO)

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19
Q

How is NH4 a dative covalent molecule

A

The ammonium ion is formed by dative covalent bonding. It forms when the nitrogen atom in an ammonia molecule donates a pair of electrons to a proton (H+)

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20
Q

What will electrons do to each other (shapes of molecules)

A

Repel each other to a point of

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20
Q

What will electrons do to each other (shapes of molecules)

A

Repel each other to a point of

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21
Q

What type of electron pair will repel the most/least

A

Lone pairs repel more than bonding pairs.

This means the greatest angles are between lone pairs of electrons, and bond angles between bonding pairs are often reduced because they are pushed together by lone pair repulsion. §1

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22
Q

What two things does shape of molecules depend on

A
  • Type of electron pairs surrounding the central atom
  • Number of electrons surrounding the central atom
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23
Q

Give the type of electron pairs and bond angles of methane, ammonia and water

A

Methane - 4bp, 0lp - 109.5
Ammonia - 3bp, 1lp - 107
Water - 2bp, 2lp - 104.5

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24
Q

Give the bond angles, electron pair configuration and name of molecules with 2 electron pairs (and examples for each)

A

Linear (BeCl2) - 2bp, 0lp - 180
Linear (double bond) (CO2) - 2bp, 2lp - 180

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25
Q

Give the bond angles, electron pair configuration and name of molecules with 3 electron pairs around central atom (and examples for each)

A

Trigonal planar (BCl3) - 3bp, 0lp - 120
Non linear (SO2) - 2bp, 1lp - 119

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26
Q

Give the bond angles, electron pair configuration and name of molecules with 4 electron pairs around central atom (and examples for each)

A

Tetrahedral (NH4+) - 4bp, 0lp - 109.5
Trigonal planar (PF3) - 3bp, 1lp - 107
Non linear (H2O)- 2bp, 2lp - 104.5

27
Q

Give the bond angles, electron pair configuration and name of molecules with 5 electron pairs around central atom (and examples for each)

A

Trigonal bipyramidal (PCl5) - 5bp, 0lp - (90 between three straight and 120 between one coming into and out of the page)
Seesaw (SF4) - 4bp, 1lp - (87 between straight one and one coming into/out of the page, 102 between the bond coming into and out of the page, there’s only one)
Distorted T (ClF3) - 3bp, 2lp - 87.5

28
Q

Give the bond angles, electron pair configuration and name of molecules with 6 electron pairs around central atom (and examples for each)

A

Octahedral (SF6) - 6bp, 0lp - 90
Square pyramidal (IF3) - 5bp, 1lp - 90x2 between ones into and out of page, 81.9x2 between ones on the page and ones come out
Square planar (XeF4) - 4bp, 2lp - 90

29
Q

are the electrostatic forces of attraction stronger in simple covalent molecules or giant covalent structure

A

Much stronger in giant covalent structure

30
Q

Give 2 examples of giant covalent structures and explain why they are able to form these giant structures

A

Carbon and Silicon can form these giant networks as they can each form four strong covalent bonds.

31
Q

Give 5 properties of giant molecular structure as a result of their strong covalent bonds

A

High melting point - need to break a lot of strong bonds before the substance melts, which takes a lot of energy

Often extremely hard - due to very strong bonds all through the lattice arrangements

Good thermal conductors - since vibrations travel easily through the stiff lattices

Insoluble - the covalent bonds mean atoms are more attracted to their neighbours in the lattice than to solvent molecules. The fact that they are all insoluble in polar solvents (like water) shows that they don’t contain ions

Can’t conduct electricity - since there are (in most giant covalent lattice structures) no charged ions or free electrons (all the bonding electrons are held in localised covalent bonds)

32
Q

Explain why (as an exception) graphite can conduct electricity

A

Carbon atoms form sheets, with each carbon atom sharing three of its outer shell electrons with three other carbon atoms. This leaves the fourth outer electron in each atom fairly free to move between the sheets, making graphite a conductor.

33
Q

what is graphene and what are its properties

A

Graphene is a sheet of carbon atoms joined together in hexagons. The sheet is just one atom thick, making it a 2D compound.

Graphene’s structure gives it some useful structures. Like graphite, it can conduct electricity as the delocalised electrons are free to move along the sheet. It’s also very strong, transparent and really light.

34
Q

What is the structure of metal elements

A

Giant metallic lattice structures

35
Q

Why do metallic lattices contain positive ions

A

the electrons in the outermost shell of the metal atoms are delocalised. this leaves a positive metal ion

36
Q

Explain metallic bonding

A

the positive metal ions are electrostatically attracted to the delocalised negative electrons. They form a lattice of closely packed positive ions in a sea of delocalised electrons.

37
Q

What is the overall metallic structure made up of

A

The overall lattice structure is made up of layers of positive metal ions, separated by layers of electrons

38
Q

Why do metals have a high melting points

A

because of the strong metallic bonding, with the number of delocalised electrons per atom affecting the melting point. The more electrons there are, the stronger the bonding will be and the higher the melting point. Mg2+ has two delocalised electrons per atom, so its got a higher melting point than Na+, which only has one. The size of the metal ion and the lattice structure also affect the melting point.

39
Q

Why are metals malleable and ductile

A

As there are no bonds holding specific ions together, and the layers of positive metal ions are separated by layers of electrons, metals are malleable (can be shaped) and are ductile (can be drawn into a wire). The layers of metal ions can slide over each other without disrupting the attraction between the positive ions and electrons

40
Q

Why are metals good thermal conductors

A

The delocalised electrons can pass kinetic energy to each other, making metals good thermal conductors.

41
Q

Why are metals good electrical conductors

A

Metals are good electrical conductors because the delocalised electrons are free to move and can carry a charge. Any impurities can dramatically reduce electrical conductivity by reducing the number of electrons that are free to move and carry charge - the electrons transfer to the impurities and form anions.

42
Q

Why are metals insoluble

A

Metals are insoluble, except in liquid metals because of the strength of the metallic bonds.

43
Q

What is electronegativity

A

the ability of an atom to attract the bonding electrons in a covalent bond

44
Q

What scale is used to measure the electronegativity of a molecule

A

the Pauling scale

45
Q

What is the most electronegative molecule and give 3 other examples of very electronegative molecules

A

Fluorine is the most electronegative

Oxygen, chlorine, nitrogen are all very electronegative

46
Q

more electronegative elements have (1) nuclear charges (the number of protons in the nucleus) and (2) atomic radii

A

1 - higher
2 - smaller

47
Q

electronegativity (1) across periods and up the groups (ignoring the noble gases)

A

increases

48
Q

Why are the covalent bonds in homonuclear diatomic gases non polar

A

the covalent bonds in homonuclear, diatomic gases (eg. H2, Cl2) are non polar because the atoms have equal electronegativities and so electrons are equally attracted to both nuclei

49
Q

What happens to when two atoms have different electronegativities (polarity)

A

If the bond is between two atoms with different electronegativities, the bonding electrons will be pulled towards the more electronegative atom. this causes the electrons to be spread unevenly, and so there will be a charge across the bond (each atom has a partial charge - one atom is slightly positive, and the other is slightly negative). The bond is polar

50
Q

What are the only molecules that can be purely covalent

A

Only bonds between atoms of a single element, like diatomic gases, can be purely covalent

51
Q

What type of bonding do all atoms and molecules form

A

London forces

52
Q

Explain the basic concept of a temporary dipole

A

Electrons in charge clouds are always moving really quickly. At any particular moment, the electrons in an atom are likely to be more to one side than the other. At this point, the atom would have a temporary (or instantaneous) dipole.

53
Q

Explain the concept of London forces

A

A temporary dipole can induce another temporary dipole in the opposite direction on a neighbouring atom. the two dipoles are then attracted to each other. The second dipole can induce can induce yet another dipole in a third atom. It’s like a domino effect
Because the electrons are constantly moving, the dipoles are being created and destroyed all the time. Even though the dipoles keep changing, the overall effect is for the atoms to be attracted to each other.

54
Q

Explain how iodine molecules are held together in a lattice

A
  • Iodine atoms are held together in pais by strong covalent bonds to form molecules of I2
  • But the molecules are then held together in a molecular lattice arrangement by weak London forces
  • This structure is known as a simple molecular structure
55
Q

What property of molecules will mean they have larger London forces?

A
  • Larger molecules have larger electron clouds, meaning stronger London forces.
  • Molecules with greater surface areas also have stronger London forces because they have a bigger exposed electron cloud
56
Q

Explain why molecules with stronger London forces will have higher boiling points. (and solids)

A

When you boil a liquid, you need to overcome the intermolecular forces, so that the particles can escape from the liquid surface. It stands to reason that you need more energy to overcome stronger intermolecular forces, so liquids with stronger London forces will have higher boiling points

(same applies with solids and melting points)

57
Q

Explain why, as alkane chains get longer, it becomes harder to separate the alkanes

A

Alkanes have covalent bonds inside the molecules. Between the molecules there are London forces, which hold them all together.
The longer the carbon chain, the stronger the London forces - because there’s more molecular surface contact and more electrons to interact
So, as the molecules get longer, it gets harder to separate them because it takes more energy to overcome the London forces.

58
Q

Explain why fewer London forces can formed on branched alkanes

A

Branched-chain alkanes can’t pack closely together and their molecular surface contact is small compared to straight chain alkanes of similar molecules mass. So, fewer London forces can form

59
Q

polar molecules have permanent (1) bonds

A

permanent dipole-dipole bonds (as a result of the slightly positive and negative charged on polar molecules which cause weak electrostatic forces of attraction between molecules)

60
Q

Molecules that can form permanent dipole-dipole bonds in addition to their London forces will generally have (1) boiling points and melting points than those with similar London forces that can’t form permanent dipole-dipole bonds

A

higher

61
Q

Hydrogen bonding only occurs when Hydrogen bonds to what 3 elements?

A

Fluorine, Nitrogen or Oxygen

62
Q

Fluorinem Nitrogen and Oxygen are very (1) and so they draw the bonding electrons (2) from the hydrogen atom

A

1 - electronegative
2 - away

63
Q

Explain why hydrogen atoms form weak bonds with lone pairs of electrons on the fluorine, nitrogen or oxygen atoms of other molecules

A

The bond is so polarised and hydrogen has such a high charge density because its so small that the hydrogen atoms form weak bonds with lone pairs of electrons on the fluorine, nitrogen or oxygen atoms of other molecules

64
Q

What groups will often form hydrogen bonding

A

Alcohols (-OH) and amines (-NH)

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