3.1 Physical Chemistry: .3 Bonding

Question 1

what are the ions involved in these compounds;

Hyrdochloric Acid

sulphuric acid

Nitric Acid

Ethanoic Acid

Carbonoic Acid

Phosphoric acid

Sodium Hydroxide

Ammonium

Answer 1

• Hydrochloric acid*= H⁺ … Cl^-
• sulphuric acid*= H⁺ H⁺ … SO₄^2-
• Nitric Acid*= H⁺ … NO₃^-
• Ethanoic Acid*= CH₃COO^- … H⁺
• Carbonoic Acid*= H⁺ H⁺ … CO₃^2-
• Phosphoric acid*= H⁺ H⁺ H⁺ … PO₄^3-
• Sodium Hydroxide*= Na⁺ … OH^-
• Ammonium*= NH₄⁺ … OH^-

Question 2

work out the formula of the compound when you have these following substances;

• Aluminimum Chloride
• Aluminium Sulphate
• Aluminium Nitrate
• Aluminium carbonate
• Aluminium Phosphate
• Aluminium Hyrdoxide

To work this out, you find out the charge on aluminium when it has become stable and the charge of the acids or hydroxide. The formula will depend on the amount of ions you need to cancel out the opposite charge.

Answer 2

• -Aluminimum Chloride*= Al³⁺ Cl^- ( AlCl₃ )
• -Aluminium Sulphate=* Al³⁺ SO₄^2- ( Al₂ (SO₄)₃ )
• -Aluminium Nitrate=* Al³⁺ NO₃^- ( Al(NO₃)₃ )
• -Aluminium carbonate=* Al³⁺ CO₃^2- ( Al₂ (CO₃)₃ )
• -Aluminium Phosphate=* Al³⁺ PO₄^3- ( AlPO₄)
• -Aluminium Hyrdoxide=* Al³⁺ OH^- ( Al (OH)₃ )

Question 3

what are the different types of strong chemical bonds?

Answer 3

ionic, covalent and metallic

Question 4

properties of an ionic bond

Answer 4

• ionic bonds occur between metals and non metals
• electrons are transferred from metal atoms to non metal atoms
• positive and negative ions are formed
• and electrostatic attraction holds positive and negative ions together

Question 5

which chemical has a latice of that of a cube shape and explain what a lattice is

Answer 5

sodium chloride

-a lattice is just a general structure

Question 6

what are the properties of ionically bonded compounds?

Answer 6

• always solid at room temp
• ionic compounds can conduct electricity when molten or dissolved not solid because the ions are free to move and they carry a charge, whlist in a solid the ions are in a fixed position by strong ionic bonds
• ionic compounds have high melting points because giant ionic lattices are helf together by strong electrostatic forces and it take high enery to overcome these forces
• Ionic compounds tend to dissolve in water because water molecules are polar. So part of the molecule has a small negative charge and the other a small positive charge. These charged parts pull ions away from the lattice causing it to dissolve,
• ionic compounds are brittle and can shatter easily when given a sharp blow because they form a lattice of alternating positive and negatice ions so a blow may move the ions and produce contact between ions with like charges

Question 7

what are electrolytes?

Answer 7

electrolytes are ionic salts (e.g Sodium Chloride)

• they conduct electricity as a molten liquid or in solution but not as a solid.
• the ions act as charged carriers

Question 8

what is covalent bond

Answer 8

a shared pair of electrons

Question 9

properties of a covalent bond

Answer 9

• a covalent bond forms between a pair of non-metal atoms
• the atoms share some of their outer electrons so that each atom has a stable noble gas arrangement
• a single covalent bond is shared pair of electrons

Question 10

in a double covalent bond, how many electrons are shared?

Answer 10

4 electrons

Question 11

how does sharing elctrons hold atoms together?

Answer 11

atoms with covalent bonds are held together by the electrostatic attraction between the nuclei and the shared electrons. This takes place within the molecule

Question 12

do covalent bonds break during melting and boiling?

Answer 12

No because during melting and boiling you only overcome intermolecular forces holding the molecules together in the covalent structures

Question 13

what is coordinate or dative bonding?

Answer 13

where both a pair of electrons comes from one atom

Question 14

show how ammonia NH3 can become NH4+

Answer 14

the nitrogen atom of ammonia uses it lone pair of elctrons to form a coodinate bond with the H+ ions of [an acid maybe].

Question 15

explain the arrangment of cabon atoms in graphite

Answer 15

the carbon atoms are arranged in sheets of flat hexagons covalenty bonded with three bonds each (carbon can make 4 bonds). The fourth outer electron of each carbon is delocalised

-the sheets of hexagns are bonded together by weak van der waals forces

Question 16

what are the properties of graphite?

Answer 16

• the weak bonds between the layers of graphite are easily broken, so the sheets can slide over each other thats why it feels slippery and can be used as dry lubricant
• it is an electrical conductor because the delocalised electron in graphite isnt attached to any particular carbon atom and so is free to move along the sheets carrying a charge
• layers are quite far apart compared to the lenght of the covalent bonds so graphite has a low density and is used to make strong light weight sports equipment
• has a high melting point as energy is needed to overcome the covalent bonds in the hexagonal sheet
• graphite is insoluble in any solvent, the covalent bonds are too strong to break.

Question 17

describe the structure of diamond

Answer 17

diamond is made up of carbon atoms

each carbon is covalently bonded to to four other carbon atoms

the atoms arrange themselves in a tetrahelal shape/

Question 18

what are the properties of diamond?

Answer 18

• diamond has very high melting point because of the strong covalent bonds that require a lot of energy to overcome
• diamond is extremely hard which makes it good for drills and saws
• it is a good thermal conductor because vibrations travel easily through the stiff lattice
• it can not conduct electricity because all outer electrons are used up in bonding
• diamond does not dissolve in any solvent
• you can cut diamond to make gem stones, its structure makes it refract light a lot

Question 19

what is a metallic bond?

Answer 19

metallic bonding involves attraction between delocalised electrons and positive ions arranged in a lattice

Question 20

what would a simple picture of metallic bonding show?

Answer 20

a metal consisting of a lattice of positive ions existing in a sea of out elerctrons which are delocalised.

Question 21

what are the properties of metals?

Answer 21

• metals are good conductors of electricity; the delocalised electrons that carry charge allow electricity to be conducted
• good conductors of heat; they have high thermal conductiviities due to the sea of delocalised electrons which can pass kinectic energy to each other
• malleable and ductile; malleable (beaten into shape), ductile (pulled into thin wires)
• metals have high metling and boiling points; because they are giant structures. There is a strong attractoin between metal ions and the delocalised sea of electrons which makes the atoms difficult to seperate. The number of delocalised electrons per atom the stronger the bond will be and the higher the melting point e.g Mg2+ has a melting point than Na⁺
• metals are insoluble because of the strong metallic bonds

Question 22

what does the strenght of a metal depend on?

Answer 22

• the charge of the ion; the greater the charge, the greater the number of delocalised electrons and the stronger the electrostatic force betwee the positiv ions and electrons
• the size of the ions; the smaller the ion, the close the electrons are to the positive nucleus and the stronger the bond.

Question 23

what are the four types of crystal structures?

Answer 23

• ionic
• metallic
• macromolecular
• molecular

Question 24

give an example of a ionic crytal ctructure

and state whether it has a high or low boiling point, what state its in at room temperature, whether it conducts electricity as a solid and a liquid and is soluble in water

Answer 24

sodium chloride (NaCl)

• it has a high melting and boiling point
• it is a solid at room temperature ]
• it does not conduct electricity when solid as there are no freely moving ions but does conduct electricity when liquid as the ions are free to move
• it is soluble in water

Question 25

why are ionic crystal structures sobule in water?

Answer 25

• many ionic solids are soluble in water- although not all.
• It depends on whether there are big enough attractions between the water molecules and the ions to overcome the attractions between the ions themselves.
• Positive ions are attracted to the lone pairs on water molecules and co-ordinate (dative covalent) bonds may form. Water molecules form hydrogen bonds with negative ions.

Question 26

give an example of a metallic crystal structure

and state whether it has a high or low boiling point, what state its in at room temperature, whether it conducts electricity as a solid and a liquid and is soluble in water

Answer 26

Magnesium (Mg)

• this crystal structure has a high boiling point
• it is a solid at room temperature
• it does conduct electricity when solid and liquid because of the sea of delocalised electrons
• it is not soluble in water

Question 27

give an example of a macromolecular (giant covalent) crystal structure

and state whether it has a high or low boiling point, what state its in at room temperature, whether it conducts electricity as a solid and a liquid and is soluble in water

Answer 27

• diamiond/ Graphite/ SiO₂
• giant covalent crystal structures have high melting and boiling points
• they are solid at room temperatures
• only graphite conducts electrecity because of the lone electron that is able to move along the layer and they can conduct electricity as sublimes rather than liquid
• they are not soluble in water

Question 28

give an example of a molecular (simple covalent) crystal structure

and state whether it has a high or low boiling point, what state its in at room temperature, whether it conducts electricity as a solid and a liquid and is soluble in water

Answer 28

• Carbon dioxide (CO₂)/ Iodine (I₂)/ Water (H₂O)
• they have a low melting and boiling point because the intermolecular forces dont need a lot of energy but the covalent bonds are still not broken
• Iodine is a solid at room temperature but the other are usually liquid or gas
• they do not conduct electricity when liquid or solid because they have no free electrons as all electrons are used in bonding and they do not have an overall electric charge.
• their solublity depends on how polarised the molecule is

Question 29

describe the structure of ice

Answer 29

• each ice molecule forms up to a regular hexagonal lattice out of water molecules which bond to each ther through hydrogen bonding
• the hydrogen bonds form between the lone pair of electrons on the oxygen and the hydrogen atom

Question 30

what are the properties of ice?

Answer 30

• ice has a high melting point compared to other simple molecular structures due to the higher amount of hydrogen bonding
• ice is also less dense than water because when water freezes there is an expansion as the hydrogen bond establishes. Each hydrogen bonds to four others with a tetrahedal arrangement about the oxygen. Hence there is a lot of empty space between molecules of water in ice.

Question 31

Why is iodine a solid at room temperature?

Answer 31

• iodine has a herring bone structure
• between each Iodine atom there is a covalent bond
• between each iodine molecule there are van der waals forces
• I₂ has a 106e^- therefore the van der waals forces is greater, melting point higher which explains why it is a solid at room temperature

Question 32

talk about the relationship between angles and lone pair and bonding pair of electrons

Answer 32

lone pair to lone pair angles are bigger than lone pair to bond paire angles which are bigger than bond pair to bond pair angles.

Question 33

state the shape and bond angle of a molecule that has 2 electron pairs (no lone pair). Give an example.

Answer 33

• shape; linear
• bond angle; 180 (the 2 bond pairs repel equally)
• example; BeCl₂

Question 34

state the shape and bond angle of a molecule that has 3 electron pairs (no lone pair). Give an example.

Answer 34

• shape; triganol planar
• bond angle; 120 (the 3 bond pairs repel repel equally)
• example: BF₃

Question 35

state the shape and bond angle of a molecule that has 4 electron pairs (no lone pair). Give an example.

Answer 35

• shape: tetrahedal
• bond angle: 109.5 (all 4 pair of electrons repel equally)
• example; NH₄ ⁺

Question 36

to predict the shape of a molecule follow these steps

Answer 36

1. ) first work out which one of the central atom (thats the one all the other atoms are bonded to)
2. ) use the periodic table to work out the number of electrons in the outer shell od the central atom
3. ) add one to this number for every atom that the central atom is bonded to
4. ) divide by 2 to find the number of electron pairs on the central atom
5. ) compare the number of the electron pairs to the number of bonds to find the number of lone pairs and the number of bonding pairs on the central atom

Question 37

state the shape and bond angle of a molecule that has 4 electron pairs (2 of which are lone pair). Give an example.

Answer 37

• shape: bent
• bond angle: 104
• example; H₂O

Question 38

state the shape and bond angles of a molecule that has 5 electron pairs (no lone pair). Give an example.

Answer 38

• shape; triganol bi-pyramid
• bond angle; 90 and 120
• example; PCl₅

Question 39

state the shape and bond angle of a molecule that has 5 electron pairs (1 of which is a lone pair). Give an example.

Answer 39

• shape; Seesaw/ Tetrahedal
• bond angle; 102/109.5 (lone pair to bond pair repulsion is greater than bond pair to bond pair repulsion)
• example; SF₄

-

Question 40

state the shape and bond angle of a molecule that has 5 electron pairs (2 of which are lone pair). Give an example.

Answer 40

• shape; T shape/ Triganol Pyramid
• bond angle; 88/ 104.5
• example; CLF₃

Question 41

state the shape and bond angle of a molecule that has 6 electron pairs (no lone pair). Give an example.

Answer 41

• shape; octahedral
• bond ange; 90 (all pair of electrons repel equally)
• example; SF₆

Question 42

state the shape and bond angle of a molecule that has 6 electron pairs (2 of which are lone pair). Give an example.

Answer 42

• shape; square planar
• bond angle; 90
• example; XeF₄

Question 43

state the shape and bond angle of a molecule that has 4 electron pairs (1 of which is a lone pair). Give an example.

Answer 43

• shape; triganol pyramid
• bond angle; 107.5
• example; PF₃

Question 44

state the shape and bond angle of a molecule that has 6 electron pairs (1 of which is a lone pair). Give an example.

Answer 44

• These are very rare
• the shape is called a square pyramidal
• bond angle: 90 degrees

Question 45

what does electronegativity mean?

Answer 45

this is the atom’s ability to attract the electron pair in a covalent bond

Question 46

what is a non polar molecule?

Answer 46

where you have a covalent bond between two atoms that are of the same (equal) electronegativity causing the pair of electrons to be equally attracted to each atoms nuclei.

-e.g this will occur in moecules such as H₂

Question 47

when does a molecule become polar?

Answer 47

when you have a covalent bond between two atoms of different electronegativies

so the shair pair of electrons is more attracted to the the more electronegative atom

Question 48

Give an example of a polar molecule and a non polar molecule and say why

Answer 48

polar molecule= H₂O

Water is polar because…

the hydrogen atoms are less electronegative than the oxygen atom therefore a dipole occurs. In the shape of the molecule you can see that the pole with the hydrogen atoms has less electron density whlist the oxygen atom pole has a greater electron density (uneven distribution of charge= dipole)

non polar molecule= CO₂

Carbon dioxide is non polar because…

carbon is more electronegative than oxygen therefore a dipole occurs, however the shape of this molecule is symetrical so the dipoles cancel each other out.

Question 49

what is a dipole? so when do you get a permanent dipole?

Answer 49

when there is a difference in charge between the two atoms caused by a shift in electron density

-when you have a polar bond, which has a difference in electronegativity, causing a permanent dipole

Question 50

what does δ+ and δ− mean?

Answer 50

• these are partial charges
• δ+ means slightly positive (lesser electron density)
• δ− means slightly negative (or greater electron density)

Question 51

the greater the difference in electronegativity between the atoms the ________ polar the bond

(fill in the missing word)

Answer 51

the greater the difference in electronegativity between the atoms the more polar the bond

Question 52

why is it that some molecules with polar bonds do not have a permanent dipole?

Answer 52

in the molecule if the polar bonds are arranged symmetrically in the molecule, then the charges cancel out and there is no permanent dipole

Question 53

the molecule CCl₄ has a tetrahedral shape. Will CCl₄ be polar?

Explain your answer

Answer 53

the molecule is not polar.

each of the C-Cl bonds are polaer but the polar bonds are arranged symmetrically all around the molecule,

so the charges cancel each other out

Question 54

why are covalent compounds non polar?

Answer 54

because the atoms have equal electronegativities, so the elecrons are equally attracted to both nuclei.

Question 55

what are the different types of intermolecular forces? put them in order of weakest to strongest

Answer 55

1. Van der Waals forces (induced dipole-dipole)
2. Permanent dipole- permanent dipole forces
3. Hydrogen bonding

Question 56

explain why hydrogen Chloride has a dipole dipole bonding

Answer 56

• dipole-dipole forcs act between moecules that have permanent dipoles
• for example, in the hydrogen chloride molecule, chlorine is a lot more elctronegative than hydrogen (because it has more protons than hydrogen to attract the bonding electrons)
• so the electrons are pulled towards the chlroine atom rather than the hydrogen atom
• the molecule therefore has a dipole and is written H^δ+ and Cl^δ−

Question 57

if you put a charged rod next to a jet of polar liquid why does the liquid move?

Answer 57

• its because polar liquids contain molecules with permanent dipoles.
• if the rod is positively charged then the slightly negative pole of the liquid will move to the rod.
• if the rod is negatively charged then the slightly positively charged pole of the liquid will move to the rod.
• the polar molecules in the liquid can turn arounf so the oppositely charged end is attacted towards the rod.

Question 58

explain how van der waals forces occur

Answer 58

• electrons in charge clouds are always moving really quickly and at any moment, the electrons in an atom are likely to be on more than one side than the other. at this moment the atom would hve a temporary dipole
• the temporary dipole can cause another temporary dipole (temporary induced dipole) in the opposite direction on an neigbouring atom. The two dipoles are then attracted to each other.
• the second dipole can cause yet another dipole in a third atom. It’s kind of like a domino effect
• because the electrons are constantly movig, the dipoles are created and destroyed all the time
• even though the dipoles keep changing, the overall effect is for the atoms to be attracted to each another.

Question 59

in Iodine molecules, what bonding do you find between the atoms and between the molecules?

Answer 59

• covalent bonds between the atoms
• van der waals betweent the molecules

Question 60

what is the relationship between size of molecules, electrons and Van der Waal forces?

Answer 60

the larger the molecules have larger electron clouds (greater electron density) there for they have stronger van der Waals forces.

Question 61

how does the shape of a molecule affect van der Waal forces?

Answer 61

long straight molecules can lie closer together than branched ones

the closer together two molecules get, the stronger the forces betweent them are

Question 62

why do larger alkanes require higher boiling points?

Answer 62

because they are larger molecule there are larger electron clouds (greater electron density), which means there are stronger van der Waals forces which require more energy to overcome.

Question 63

when do hydrogen bonds occur?

Answer 63

hydrogen bonding only happens when hydrogen is covalently bonded to flourine, nitrogne or oxygen

Question 64

what functional groups do molecules have that contain hydrogen bonding?

Answer 64

-OH and -NH groups

Question 65

how do hydrogen bonds form?

Answer 65

• flourine, nitrogen and oxygen are very electronegative,
• so they draw bonding electrons away from the hydrogen atom.
• The bond is so polarised (hydrogen with its delta pluse and F, N, O with delta minus) and hydrogen has such a high charge density (because its so small)
• that the hydrogen forms weak bonds with the lone pair of electrons on the flourine, nitrogen or oxygen atoms of other molecules

Question 66

Why are the boiling points of the hydrides ( H₂O, HF, NH₃) higher than the other hyrdrides even though youd expect them to be smaller?

Answer 66

• In all hydrides there are van der waal forces operating
• but with the hydrides where oxygen, flourine and nitrogen are present there is hydrogen bonding between each molecule
• these stronger intermolecular forces of attraction make the molecules more difficult to seperate
• therefore more energy is needed

Question 67

draw a labelled diafram to show the intermolecular forces that exist between two molecules of water.

include all loned pairs and partial charges in your diagram

Answer 67