UNIT 1: SECTION 3-BONDING

Question 1

A) When do “compounds” form?

B) Outline the 2 main types of bonding in compounds

Answer 1

A)-when elements join or bond together

B)-ionic and covalent.

Question 2

A) When do ions form?

B) Describe the “simplest” type of ions?

Answer 2

A)-when 1 or more electrons transferred from 1 atom to another
B)-are single atoms which have either lost/gained electrons to get full outer shell

Question 3

C) EXAMPLES:

Briefly describe the formation of the follwing ions including their equation:

1- Na+
2- Mg2+
3- Cl-
4- O2-

Answer 3

C)1-Na atom loses 1 electron to form sodium ion
-Na–>Na+ + e-
2-Mg atom loses 2 electrons to form a magnesium ion
-Mg–>Mg2+ + e-
3-Cl atom gains 1 electron to form a chloride ion
-Cl + e- –> Cl-
4-O atom gains 2 electrons to form an oxide ion
-O + 2e- –>O2-

Question 4

D) How may you work out what ion each element forms from the periodic table?

E) Describe how ionic bonding is linked to the charges on ions

Answer 4

D)-elements in same group all have same n. of outer electrons
–>so they have to lose/gain same n. of electrons to have full outer shell aiming for
–>means will form ions with same charges
E) -electrostatic attraction holds (+) and (-) ions together–>v. strong
–>this is ionic bonding.

Question 5

A) What are “compound ions”?

B) Outline the formulas for the following ions: 
1-Sulfate
2-Hydroxide
3-Nitrate
4-Carbonate
5-Ammonium

Answer 5

A)-lots of ions that are made up of groups of atoms with an overall charge
B)1-(SO4)2-
2-(OH)-
3-(NO3)-
4-(CO3)2-
5-(NH4)+

Question 6

A) How may you work out the formula of an ionic compound?

Answer 6

A)-ionic compounds made up of a (+) charged part AND a (-) charged part

• ->overall charge of any compound is 0
• ->so all negative charges in the compound must balance all (+) charges
• ->so you can use charges on individual ions present to work out formula of an ionic compound.

Question 7

EXAMPLE: A) Work out the formula of Sodium Nitrate

B) Similarly, determine the formula of Magnesium Chloride

Answer 7

A)-this compound contains Na+ (1+) and NO3- (1-) ions
–>these charges are balanced with one of each ion
–>so the formula is NaNO3
B)-it contains Mg2+ (2+) and Cl- (1-) ions
–>as a chloride ion only has a 1- charge we will need 2 of them to balance out the 2+ charge of a magnesium ion
–>so formula is MgCl2.

Question 8

A) Describe a “giant ionic lattice”

B) Outline the structure of Sodium Chloride

C) What determines the physical properties of ionic compounds?

Answer 8

A)-ionic crystals are giant lattices of ions
-a lattice is just a regular structure
-structures giant as its made up of same basic unit repeated over and over
B)-the Na+ and Cl- ions are packed together
–>the sodium chloride lattice is cube shaped
–>different ionic compounds have different shaped structures BUT they’re all still giant lattices.
C)-the structure of the ionic compounds.

Question 9

Explain the following PROPERTIES of ionic compounds:

A) Conduct electricity when molten or dissolved BUT not when solid

B) Have high melting points

C) Tend to dissolve in water

Answer 9

A)-this as ions in liquid free to move (and they carry charge)
–>in a solid the ions fixed in position by strong ionic bonds
B)-giant ionic lattices held together by strong electrostatic forces
–>takes lots of energy to overcome these forces, so melting points are v. high (e.g: 801 degrees C for NaCl)
C)-water molecules are polar–>part of molecule has small (-) charge and other bits have small (+) charges
–>these charged parts pull ion away from the lattice
–>causing it to dissolve.

Question 10

A) Describe how molecules form

B) How do molecules bond to each other?

C) State some examples of molecules?

Answer 10

A)-molecules form when 2+ atoms bond together–> doesn’t matter if atoms same or different
B)-molecules held together via strong covalent bonds
–>single covalent bond contains shared electron pair
C)-chlorine gas (Cl2)/Carbon monoxide (CO)/ Water (H20)/Ethanol (C2H5OH)

Question 11

D) Briefly explain how covalent bonding works

E) Outline some examples of covalently bonded molecules

Answer 11

D)-2 atoms share electrons so they’ve both got full outer shells of electrons
–>both the (+) nuclei are attracted electrostatically to the shared electrons
E)-iodine: shares 1 electron pair
-hydrogen chloride: shares an electron pair
-hydrogen: shares 1 electron pair
-water (H2O) share 2 electron pairs
-methane (CH4) shares 4 electron pairs.

Question 12

A) What are “double” and “triple” bonds that may form in covalent bonding?

A) Describe “giant covalent structures”

Answer 12

A)-atoms don’t just form single bonds–>double or even triple covalent bonds can be formed between atoms too

• ->these multiple bond contain multiple shared electron pairs.
• ->E.G: CO2/N2

A)-these giant covalent structures have huge network of covalently bonded atoms
–>sometimes called macromolecular structures

Question 13

B) Why can carbon form giant covalent structures and what are the main two?

Answer 13

B)-carbon atoms may form these giant covalent structures as each can form 4 strong covalent
-2 main structures are graphite and diamond

Question 14

A) Briefly explain how each of the following structural features of “graphite” relate to it’s properties:

1-weak bonds between layers
2-“delocalised” electrons
3-layers fairly far apart compared to covalent bond length
4-strong covalent bonds in hexagon sheets
5-insoluble in solvents

B) How are the sheets of hexagons bonded together?

Answer 14

A)1-these weak bonds easily broken so sheets able to slide over each other
–>so graphite feels slippery + is used as dry lubricant + in pencils
2-delocalised electrons not attached to any particular carbon atoms + are free to move along sheets carrying a charge–>therefore is electrical conductor
3-so graphite has low density + is used to make strong, lightweight sports equipment
4-therefore graphite has v. high melting point (sublimes at over 3900 K)
5-this is as covalent bonds in sheets are too strong to break.
B)-by weak van der Waals forces.

Question 15

A) What is the basic structure of “diamond”?

Answer 15

A)-its also made up of carbon atoms

• ->each C atom covalently bonded to 4 other C atoms
• ->atoms arrange themselves in a tetrahedral shape

Question 16

B) Due to the strong covalent bonds what main features does diamond have?

Answer 16

B)1-v. high melting point–>also sublimes at 3900 K +
2-v. hard–>its used in diamond-tipped drills/saws
3-vibrations easily travel through stiff lattice so is good thermal conductor
4-can’t conduct electricity–>all outer electrons held in localised bonds
5-like graphite (same reason) diamond not dissolve in any solvent
6-can “cut” diamond to form gemstones
–>it’s structure makes it refract light a lot–>this why it sparkles.

Question 17

A) Briefly explain how an ammonium ion (NH4+) is formed and an example of what type of bonding it’s formed from

Answer 17

A)-it’s an example of dative covalent (or co-ordinate) bonding
–>it forms when the nitrogen atom in an ammonia molecule donates a pair of electrons to a proton (H+)

Question 18

A) In which form do molecules and ion come?

B) What does the shape of these ions/molecules depend on?

C) EXAMPLE: How many bonding and lone pairs of electrons does ammonia have?

Answer 18

A)-come in lots of different shapes
B)-depends on the n. of pairs of electrons in outer shell of central atom
C)-in ammonia outermost shell of nitrogen atom contains 4 pairs of electrons
–>1 lone pair (non-shared) AND 3 bonding pairs (shared pairs).

Question 19

A) How do bonding and lone pairs of electrons exist?

B) Define a “charge cloud”

Answer 19

A)-as charge clouds
B) -area where have big chance of finding an electron pair
–>electrons don’t stay still they whizz around inside the charge cloud.

Question 20

A) What is the effect of electrons being (-) charged on charge clouds?

B) What is the effect of the shape of the charge cloud on how much they repel each other (the “valence-shell Electron-Pair Repulsion Theory”)

Answer 20

A)-charge clouds will repel each other as much as they can
–>so electron pairs in outer shell of atom will sit as far apart from each other as possible
B)-lone-pair charge clouds repel more than bonding-pair charge clouds
–>so greatest angles are between lone pairs of electrons (lone-pair/lone-pair)
–>and bond angles between bonding pairs are often reduced as they are pushed together by lone-pair repulsion. (lone-pair/bonding-pair)
-smallest angles are between (bonding-pair/bonding-pair)

Question 21

A) Outline the steps that you must follow to predict the shape of a molecule using the n. of electron pairs

Answer 21

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

Question 22

For each of the following electron pairs outline the SHAPE + ANGLE(S) + EXAMPLE of each:

A) 2 electron pairs 
B) 3 electron pairs
C) 4 electron pairs 
D) 5 electron pairs 
E) 6 electron pairs

Answer 22

A)-no lone pairs/ linear shape/ 180 degrees/ BeCl2
B)-no lone pairs/ trigonal planar shape/ 120 degrees/ BF3
C)-no lone pairs: tetrahedral shape/ 109.5 degrees/ NH4+
–>1 lone pair: trigonal pyramidal shape/ 107 degrees/ PF3
–>2 lone pairs: bent shape/ 104.5 degrees/H2O
D)-no lone pairs: trigonal bipyramidal/ 90 and 120 degrees/ PCl5
-1 lone pair: seesaw shape/ 87 and 102 degrees/ SF4
-2 lone pairs: T-shaped/ 88 degrees/ ClF3
E)-no lone pairs: octahedral shape/ 90 and 90 degrees/SF6
-2 lone pairs: square planar shape/ 90 degrees/ XeF4.

Question 23

A) EXAMPLE: Predict the shape of the molecule H2S

Answer 23

A)1-central atom is sulfur
2- Sulfur in group 6 so it has 6 electrons in it’s outer shell to start with
3-Sulfur atom bonded to 2 hydrogen atoms so it has (6+2)= 8 electrons in it’s outer shell in H2S
4-n. of electron pairs on central sulfur atom is 8/2=4 pairs
5-Sulfur atom has 4 electron pairs and has made 2 bonds–>so it has 2 bonding pairs + 2 lone pairs
–>this means that H2S will have bent shape like water

Question 24

A) Define “electronegativity”

B) What is the most electronegative element and which other elements after it are also fairly electronegative?

C) State the electronegativity of H, C, N, Cl, O and F (electronegativity on the pauling scale)

Answer 24

A)-an atoms ability to attract the electron pair in a covalent bond
B)-fluorine is most electronegative
–> oxygen/nitrogen/chlorine are also strongly electronegative.
C)-H=2.2/ C=2.5/ N=3.0/ Cl=3.0/ O=3.4/ F=4.0

Question 25

A) Outline what makes a bond "polar"

Answer 25

A)-in a covalent bond between 2 atoms of different electronegativities, the bonding electrons will be pulled towards more electronegative atom-->makes bond polar.

Question 26

B) In which 2 circumstances is a bond "non-polar"?

Answer 26

B)1-when there is covalent bond between 2 atoms of same element (e.g: H2) as atoms have equal electronegativities so electrons equally attracted to both nuclei 2-some elements like carbon and hydrogen have pretty similar electronegativities so bonds between them essentially non-polar.

Question 27

C) Explain what a "permanent dipole" is and how it's created D) What affects the strength of a polar bond?

Answer 27

C)-a dipole is a difference in charge between 2 atoms that causes a shift in electron density in the bond -in polar bond, electronegativity difference between 2 atoms causes permanent dipole D)-greater the difference in electronegativity between the atoms-->more polar the bond is

Question 28

E) EXAMPLE: Why does HCL have a permanent dipole?

Answer 28

E)-chlorine much more electronegative than hydrogen, so hydrogen chloride has a permanent dipole - -> H is slightly (+)-->delta + - -> Cl is slightly (-)-->delta -

Question 29

A) When is a molecule "polar"? B) When may a molecule not be "polar"?

Answer 29

A)-happens when have molecule that contains polar bonds-->you may end up with uneven distribution of charge across whole molecule-->molecule ends up polar B)-not all molecules contain polar bonds -->if polar bonds are arranged symmetrically in molecule-->then the charges cancel out and there is no permanent dipole.

Question 30

A) How may permanent-dipole forces arise between polar molecules?

Answer 30

A)-in substance made up of molecules that have permanent dipoles, there will be weak electrostatic forces of attraction between the delta + and delta - charges on neighbouring molecules.

Question 31

B) Why does hydrogen chloride gas have polar molecules?

Answer 31

B)-the delta - chlorine is attracted to the delta + hydrogen on the next HCl molecule.

Question 32

C) If you put a charged rod next to a jet of a polar liquid like water why will the liquid move towards the rod?

Answer 32

C)-this is as polar liquids contain molecules with permanent dipoles - ->it doesn't matter if rod is (+) or (-) charged - ->polar molecules in liquid able to turn around so that oppositely charged end attracted towards rod

Question 33

A) Define "intermolecular forces" B) What are the 3 types of intermolecular forces? C) Why are intermolecular forces important?

Answer 33

A)-are forces between molecules -->are much weaker than covalent/ionic/metallic bonds B)-induced dipole-dipole OR Van der waals forces -permanent dipole-dipole forces (caused by polar molecules) -hydrogen bonding. C)-as they affect physical properties of a compound.

Question 34

A) Explain how "Van der Waals forces" work

Answer 34

A)1-electrons in charge cloud always moving quickly -->at any moment, electrons in atom more likely to be on one side than other -->at this moment atom would have a temporary dipole 2-this dipole can cause yet another temporary dipole in opposite direction on neighbouring atom-->2 dipoles then attracted to each other 3- 2nd dipole can then cause yet another dipole in 3rd atom (domino type of effect) 4-as electrons constantly moving dipoles are being created + destroyed all time -->even though dipoles keep changing the overall effect is for atoms to be attracted towards each other.

Question 35

A) How can van der waals forces hold molecules in a lattice (example-Iodine, I2)?

Answer 35

A)-Iodine (I2) is solid at room temp - ->the van der waals forces between iodine molecules that are responsible for holding them together in lattice - 1: iodine atoms held together in pairs by strong covalent bonds to form I2 molecules - 2: but the molecules held together in molecular lattice arrangement via weak van der waals attractions.

Question 36

A) Outline how the size of the molecule affects Van der Waals forces B) How does a shape's molecule affect the strength of Van Der waal's forces?

Answer 36

A)-larger molecules have larger electron clouds-->so stronger van der waal's forces B)-long, straight molecules lie closer together than branched ones -->closer together 2 molecules get the stronger forces between them are

Question 37

C) Describe how the strength of van der waals forces affect their boiling point

Answer 37

C)-when boil liquid need to overcome intermolecular forces so that the particles can escape from liquid surface - ->stands to reason that need more energy to overcome stronger intermolecular forces - -> so liquids with stronger van der waals forces have higher boiling points.

Question 38

D) EXAMPLE: Explain what happens to the boiling point of straight chain alkanes as their length increases?

Answer 38

D)-as alkane chain longer the n. of electrons in molecule increases -->means van der waals forces stronger and so the boiling points increase.

Question 39

A) Describe "hydrogen bonding"

Answer 39

A)-fluorine/nitrogen/oxygen are v. electronegative so they draw bonding electrons away from hydrogen atom -->bond is so polarised + hydrogen has such high charge density (as it's so small) that hydrogen atoms form weak bonds with lone pairs of electrons on fluorine/nitrogen or oxygen atoms of other molecules

Question 40

B) When does hydrogen bonding occur?

Answer 40

B)-only happens when hydrogen covalently bonded to fluorine/nitrogen or oxygen.

Question 41

C) Which groups do molecules that contain hydrogen bonding usually have? D) How does hydrogen bonding occur between water (H2O) molecules and ammonia (NH3) molecules?

Answer 41

C)-contain (-OH) or (-NH) groups D)-H2O: a lone pair of electrons on oxygen is attracted to the hydrogen -NH3: a lone pair of electrons on nitrogen is attracted to the hydrogen

Question 42

Hydrogen bonding's effect on the properties of substances: E) Why do substances with hydrogen bonds have higher boiling + melting points than other similar molecules?

Answer 42

E)-this is due to extra energy needed to break hydrogen bonds -->this is case with water + hydrogen fluoride which has much higher boiling point than other hydrogen halides

Question 43

F) Explain why ice is less dense than liquid water

Answer 43

F)-as liquid water cools to form ice-->molecules make more hydrogen bonds + arrange themselves into regular lattice structure - ->in this regular structure the H2O molecules are further apart on average than liquid water molecules - ->therefore ice less dense than liquid water.

Question 44

A) In what form do most metal elements exist? B) Explain "metallic bonding"

Answer 44

A)-as giant metallic lattice structures B)-outermost shell electrons of a metal atom is delocalised-->electrons free to move about the metal -->leaves a (+) metal ion. E.G: Na+/Mg2+ etc.) -the (+) metal ions are attracted to delocalised (-) electrons -->they form a lattice of closley packed (+) ions in sea of delocalised electrons -->this is metallic bonding

Question 45

A) Explain why metals have HIGH MELTING POINTS B) How does the n. of delocalised electrons per atom affect a metal's boilinig point?

Answer 45

A)-this is due to strong electrostatic attraction between (+) metal ions and the sea of delocalised electrons B)-more delocalised electrons there are the stronger the bonding-->so melting point higher -->E.G: Mg2+ has 2 delocalised electrons per atom so it's boiling point higher than Na+ which only have 1

Question 46

C) Outline why metals are good THERMAL CONDUCTORS D) What make metals good ELECTRICAL CONDUCTORS? E) Why are metals INSOLUBLE?

Answer 46

C)-the delocalised electrons able to pass K. Energy to each other D)-this is as the delocalised electrons are able to move and carry charge. E)-this is due to the high strength of the metallic bonds (except in liquid metals).

Question 47

Describe the behavior of the particles in each of the following: A) SOLID B) LIQUID

Answer 47

A)-particle are v. close together -->this gives it high density + makes it incompressible -->particles vibrate about a fixed position + can't move about freely B)-typical liquid density similar to a solid + virtually incompressible -->particles move about freely + randomly within the liquid-->allowing it to flow

Question 48

C) GAS D) What needs to be done in order to change from a solid to liquid to gas?

Answer 48

C)-gas particles have lot's energy and are further apart -->so density is generally pretty low + it's v. incompressible -->particles move about freely with not much attraction between them -->so they diffuse to fill container D)-need to break forces holding particles together -->to do this need to give particles energy (e.g: heating them).

Question 49

A) Explain why simple convalent compounds have relatively low melting/boiling points

Answer 49

A)-only have to overcome intermolecular forces that hold the molecules together -->don't need to break much stronger covalent bonds that hold the atoms together in the molecules

Question 50

B) EXAMPLE: Compare and contrast the melting and boiling points of Chlorine (Cl2) and Pentane (C5H12) with Diamond

Answer 50

B)1-Cl2: is simple covalent substance with M.P of -101 degrees C AND B.P of -34 degrees C - ->so is gas at room temp + pressure - C5H12 also a simple covalent compound has M.P -130 degrees C AND B.P of 36 degrees C - ->so is liquid at room temp + pressure 2-DIAMOND: by contrast is giant covalent substance - ->so covalent bonds do need to be broken between atoms to turn it in to liquid OR gas - ->never really melts but sublimes at over 3600 degrees C.

Question 51

Describe the way as to how particles make up substances affect the following properties?: A) Melting and Boiling points B) Conductivity of electricity C) Solubility of a substance

Answer 51

A)-these are determined by strength of attraction between substances particles B)-substance will only conduct electricity if contains charged particles free to move C)-how soluble substance is in water depends on type of particles that it contains -->water is polar solvent so polar substances OR charged ones will dissolve well in it -->BUT non-polar OR uncharged substances wont.

Question 52

Outline the nature of the following properties for each of the different types of bonding: ``` 1-Examples 2-M.P and B.P 3-typical state at room temp + pressure 4-Does Solid conduct electricity? 5-Does liquid conduct electricity? 6-Is it soluble in water? ``` A) IONIC B) SIMPLE COVALENT (MOLECULAR) C) GIANT COVALENT (MACROMOLECULAR) D) METALLIC

Answer 52

A)1-NaCl/ MgCl2 2-High 3-Solid 4-No (ions held in place) B)1-CO2/ I2/ H2O/ 2-low (inovolves breaking intermolecular forces but not covalent bonds) 3-may be solid like I2 but usually liquid OR gas 4-No 5-No 6-depends on how polarised molecule is C)1- Diamond/ Graphite/ SiO2 2-High 3-Solid 4-No (except graphite) 5-sublimes rather than melts 6-No D)1- Fe/Mg/Al 2-High 3-Solid 4- Yes (delocalised electrons) 5-Yes (delocalised electrons) 6-No.

Question 53

You can use the properties of a material to predict it's structure: A) EXAMPLE: Substance X has a melting point of 1045 K. When solid, it is an insulator, but once melted it conducts electricity. Identify the type of structure present in substance X.

Answer 53

A)1-substance X not conduct electricity when solid BUT does once melted -->so could be ionic as fits fact that it has high melting point too 2-it's definitely not simple covalent as has high melting point AND definitely not metallic as doesn't conduct electricity when solid AND definitely not giant covalent as does conduct electricity when melted --> so substance X must be IONIC.