UNIT 1: SECTION 1-ATOMIC STRUCTURE

Question 1

A) What are all elements made of?

B) Outline the 3 types of particle that atoms are made up of

Answer 1

A)-atoms

B)-protons/neutrons/electrons

Question 2

C) Describe “electrons” that are part of the atom

D) Briefly explain the structure of the nucleus in an atom

Answer 2

C)-have -1 charge
-they whizz round nucleus in orbitals which take up most volume of atom
D)-most atom mass concentrated in nucleus
-nucleus diameter smaller compared to whole atom
-nucleus is where protons + neutrons found

Question 3

E) State the RELATIVE MASS (RM) and RELATIVE CHARGE (RC) of the 3 sub-atomic particles#

F) Why is the “relative mass” and “relative charge” used for the subatomic particles?

Answer 3

E)-proton: RM=1 / RC= +1
-neutron: RM= 1 / RC= 0
-electron, e- RM=1/2000 / RC= -1.
F)-this is as the mass/charge of these subatomic particles is v. small.

Question 4

A) What are the 3 key parts of NUCLEAR SYMBOLS

B) What may be figured out from the nuclear symbol of an element?

Answer 4

A)1- MASS NUMBER (TOP NUMBER)-tells you total protons + neutrons in nucleus
2- ATOMIC (PROTON) NUMBER-this is total protons in nucleus–> it identifies element
-all atoms of same element have same n. of protons
B)-n. of protons/neutrons/electrons

Question 5

C) When are atoms “neutral” and have no overall charge?

D) How could you calculate the n. of neutrons from the nuclear symbol of an element?

Answer 5

C)-when n. of electrons same as n. of protons

D)-when minus mass number from atomic number (i.e: top minus bottom)

Question 6

A) What are ions?

B) How do atoms form ions?

C) Outline the subatomic particles in “positive” and “negative” ions

Answer 6

A) -they have different n. of protons and electrons
B)-by losing or gaining electrons
C)-negative ions have more electrons than protons
-positive ions have fewer electrons than protons

Question 7

D) Describe a Br- ion and an Mg 2+ ion in terms of electrons

Answer 7

D)-Br-: negative charge means 1 more electron than protons

• ->Br has 35 protons so Br- must have 36 elctrons
• ->overall charge is +35 -36= -1
• Mg 2+: 2+ charge means 2 fewer electrons than there are protons
• ->Mg has 12 protons so Mg 2+ must have 10 electrons
• ->overall charge= +12-10= +2

Question 8

A) Define the term “isotope”

B) EXAMPLE: Outline how 35 Cl is different to 37 Cl

Answer 8

A)-isotopes of an element are atoms with same n.of protons but different n. of neutrons
B)-35 Cl: 35-17=18 neutrons
-37 Cl: 37-17=20 neutrons
–>different mass number means different n.of neutrons
–>atomic numbers are same as both isotopes have 17 protons and 17 electrons.

Question 9

C) Outline what determines the “chemical properties” of an element are and what they are like for different isotopes

D) What determines the “physical properties” of an element and what are the physical properties for different elements?

Answer 9

C)-its number + arrangement of electrons that decides chemical properties of an element
–>isotopes have same electron configuration so they have same chemical properties
D)-physical properties tend to depend on mass of an atom
–>isotopes of an element do have slightly different physical properties though due to different mass numbers (like densities/diffusion rate etc.)

Question 10

F) State the % composition of each of the following isotopes of magnesium in magnesium and outline the n. of protons, neutrons and electrons in each:

1-24 Mg
2-25 Mg
3-26 Mg

Answer 10

1-79%/ 12 P / 12 N/ 12 E
2-10%/ 12 P/ 13 N / 12 E
3- 11% / 12 P / 14 N/ 12 E.

Question 11

THE ATOMIC MODEL:

A) Why did the atomic model change over time?

Answer 11

A) -current model is the accepted one

• but in past completely different models accepted as they fitted evidence of time
• ->as scientists did more experiments new evidence found and models modified to fit it.

Question 12

B) How did JOHN DALTON describe the atom as at the start of the 19th Century?

C) In 1897 what was J.J THOMPSON’S input on the atomic model?

Answer 12

B)-as solid spheres and said that different spheres made up different elements
C)-he discovered electrons
–>this showed atoms not solid + indivisible
–>solid sphere idea had to be changed so new model known as “plum pudding model”.

Question 13

D) In 1909 Ernest Rutherford and his students Hans Geiger + Ernest Marsden conducted the famous “gold foil experiment”. Explain this experiment

Answer 13

D) 1-fired (+) charged alpha particles at v. thin gold sheet

• plum pudding model suggested most alpha particles would be slightly deflected by (+) pudding that made up most of atom
• ->BUT most particles passed straight through gold with small number deflected backwards

Question 14

E) What model did Rutherford conclude to from the experimental results?

Answer 14

E)-plum pudding model couldn’t be right so he developed “nuclear model” of atom
–>in this a tiny (+) charged nucleus surrounded by “cloud” of (-) electrons–>most of atom empty space.

Question 15

F) What did Niels Bohr realise about Rutherford’s nuclear model of the atom?

Answer 15

F)-if electrons were in “cloud” around nucleus of atom they would quickly spiral down into nucleus causing atom to collapse

Question 16

G) Describe the model of the atom that Bohr proposed

Answer 16

G)-proposed new atomic model where electrons exist in shells or orbits of fixed energy

• ->when electrons move between shells, elctromagnetic radiation (with fixed energy or frequency) emitted OR absorbed
• ->Bohr’s model fitted experimental observations or the radiation emitted + absorbed by atoms.

Question 17

H) What did scientists later realise was wrong with Bohr’s model?

I) Therefore, how was Bohr’s model refined?

Answer 17

H)-that not all electrons in a shell have same energy
–>meant Bohr’s model not quite right
I)-it was refined to include sub-shells
–>this refined model not quite perfect as other more accurate models exist BUT this refined model is useful as it’s simple + explains many experimental observations like bonding + ionisation energy trends.

Question 18

A) Describe what it means by “relative mass”

Answer 18

A) -actual mass of atom v. small–>far too small to weigh

–>so mass of one atom is compared to mass of different atom

Question 19

B) Define the following terms:

1-relative atomic mass, Ar
2-relative isotopic mass
3-relative molecular mass, Mr

Answer 19

B) 1-is average mass of atom of an element on scale where atom of carbon-12 is 12
2-is mass of an atom of an isotope of an element in scale where atom of carbon-12 is 12
3-is average mass of a molecule on scale where atom of carbon 12 is 12.

Question 20

C) Is relative atomic mas and relative isotopic mass a whole number?

D) How do you find the relative molecular mass?

Answer 20

C)-relative atomic mass is average so not usually whole number
-relative isotopic mass is usually whole number
D)-just add up all relative atomic mass values of all atoms in the molecule

Question 21

E) What do you do to find the relative formula mass and what is it used instead of?

Answer 21

E)-relative formula mass used instead for compounds that are ionic (or giant covalent)
–>to find relative formula mass just add up the relative atomic masses (Ar) of all atoms in the formula unit.

Question 22

A) What can “mass spectrometry” be used to find?

Answer 22

A)-relative atomic mass/relative molecular mass/relative isotopic abundance

Question 23

FUNCTIONING OF TIME OF FLIGHT MASS SPECTROMETER (TOF):
Explain what happens at each step when a sample is put in to a time of flight (TOF) mass spectrometer:

A) IONISATION- explain “electrospray ionisation” and “electron impact ionisation”

Answer 23

A)-there are 2 ways to ionise your sample
1-electrospray ionisation: sample dissolved + pushed through small nozzle at high pressure
–>high voltage applied to it–>causing each particle to gain an H+ ion
–>sample turns to gas made up of (+) ions
2-Electron impact ionisation: sample vaporised and an “electron gun” used to fire high energy electrons at it
–>knocks 1 electron off each particle so they become +1 ions.

Question 24

B) Describe the DETECTION part of mass spectrometry

C) What is ION DRIFT in mass spectrometry?

Answer 24

B)- (+) charged ions accelerated by electric field so they all have same K.E
–>means that lighter ions end up moving faster than heavier ones
C)-ions enter region with no electric field so they just drift through it
–>lighter ions will drift through faster than heavier ones.

Question 25

D) Explain the DETECTION stage mass spectrometry

Answer 25

D)-as lighter ions travel at higher speeds in drift region - ->they reach detector in less time than heavier ones - ->detector detects charged particles-->mass spectrum produced.

Question 26

A) What is a "mass spectrum"? B) On a mass/charge spectrum graph what do the following show (with reference to the example graph shown): 1 Y-AXIS 2 X-AXIS 3 EACH LINE

Answer 26

A)-is the mass/charge plotted against abundance B)1-gives abundance of ions often as % -->for element height of each peak gives relative isotopic abundance (e.g: 92.4% of this sample made up of 7 Li isotope) 2-the units are given as "mass/charge" ratio (sometimes m/z as short) 3-if sample is an element each line will represent a different isotope of the element.

Question 27

C) Still with reference to the example mass spectrum given above: Why is the mass/charge ratio of each peak the same as the relative mass of that isotope? D) If electron spray ionisation had been used instead what results would be expected from the mass/charge ratio?

Answer 27

C)-this as the spectrum above was produced using electron impact ionisation -->1 electron knocked off each particle to turn them in to +1 ions. D)-if electron spray ionisation was used instead an H+ ion would have been added to each particle to form +1 ions --so mass/charge ratio of each peak would of been 1 unit greater than the relative mass of each isotope. (all spectra shown in this topic have been produced using electron impact ionisation)

Question 28

Using the example given to calculate the relative atomic mass (Ar) of an element from the mass spectrum: A) Calculate the Ar for Magnesium using the mass spectrum

Answer 28

A)1-for each peak read % relative isotopic abundance from y-axis and relative isotopic mass from the x-axis 2-multiply them together to get the total mass for each isotope and then add them up: --> (79*24=1896) + (10*25=250) + (11*26=286)= 2432 3-divide by 100 (as % used) -->Ar (Mg)=2432/100=24.32=24.3 (3 s.f).

Question 29

B) (With reference to the mass spectrum graph given) | What would you do if the relative abundances are not given as a % so then the total abundances may not add up to 100?

Answer 29

B)-do steps 1 and 2 as above but then divide by sum of relative abundances instead of 100 -Ar (Ne)=(114*20)+(0.2 *21) +(11.2*22)/(114 + 0.2 + 11.2)=20.2 (3 s.f)

Question 30

A) How may mass spectrometry be used to identify different elements? B) If an element only has 1 stable isotope like many do how may they still be identified? (example of magnesium isotope given)

Answer 30

A)-elements with different isotopes produce more than 1 line mass spectrum as the isotopes have different masses -->this produces characteristic patterns that can be used as "fingerprints" to identify certain elements. B)-they may still be identified in a mass spectrum by looking for a line at their relative atomic mass.

Question 31

A) How may you get a mass spectrum for a molecular sample and then use it?

Answer 31

A)-molecular ion (M+) formed in mass spectrometer when 1 electron removed from the molecule - ->this gives peak in spectrum with mass/charge ratio equal to relative molecular mass of molecule - ->this may be used to help identify an unknown compound

Question 32

B) EXAMPLE (use the table given) : A sample of a straight-chain alcohol is analysed in a mass spectrometer. The mass/charge ratio of its molecular ion is 46.0. Identify the alcohol C) What results should you expect if you have a mixture of compounds with different Mr values?

Answer 32

B)-table on right shows the Mr of first 3 straight-chain alcohols -->the mass/charge ratio of the molecule ion must equal to the Mr of the alcohol in the sample --> so the alcohol must be ethanol (C2H5OH). C)-you will get peak for the molecular ion for each one.

Question 33

A) Describe electrons in terms of "energy levels" and "energies" B) What is "principal quantum number" and what affects it?

Answer 33

A)-in currently accepted atomic model electrons have fixed energies -also electrons move around the nucleus in certain regions of the atom called shells/energy levels. B)-each shell given number called the principal quantum number -->further a shell is from nucleus-->the higher it's energy and the larger it's principal quantum number.

Question 34

C) What do experiments show about the amount of energy that electrons have in a shell?

Answer 34

C)-not all electrons have exactly the same energy - ->atomic model explains this-->shells are divided up into sub-shells that have slightly different energies - ->sub-shells have different n. of orbitals which can hold up to 2 electrons.

Question 35

D) For each of the following sub-shells outline the n. of orbitals it holds and the max n. of electrons 1-s 2-p 3-d 4-f

Answer 35

D)1-s: 1 orbital--> 1*2=2 electrons (max) 2-p: 3 orbitals--> 3*2=6 electrons (max) 3-d: 5 orbitals--> 5*2=10 electrons (max) 4-f: 7 orbitals--> 7*2= 14 electrons (max)

Question 36

E) For the first 4 shells outline the sub-shells in each and the total n. of electrons in each shell F) In what direction do the 2 electrons in each orbital spin?

Answer 36

E) 1st: 1s--> 2 electrons 2nd: 2s, sp--> (2) + (3*2)=8 electrons 3rd: 3s, 3p, 3d--> (2) +(3*2) + (5*2)=18 electrons 4th: 4s, 4p, 4d, 4f--> (2) + (3*2) + (5*2) + (7*2)= 32 electrons. F)-opposite directions.

Question 37

A) State the 3 main rules you must follow to figure out electron configurations

Answer 37

A) 1-electrons fill up lowest energy sub-shells first 2-electrons fill orbitals singly before they start sharing -->Nitrogen: 1s 2, 2s 2, 3s, 3 (1 electron per orbital first so all occupied singly) -->Oxygen: 1s 2, 2s 2, 2p 4 (1st orbital in 2p full but next two have electrons in singly). 3-for configuration of ions from S and P blocks of periodic table just remove OR add electrons to or from highest energy occupied sub-shell -E.G: Mg= 1s2, 2s2, 2p6, 3s2-->so Mg += 1s2, 2s2, 2p6 -Cl= 1s2, 2s2, 2p6, 3s2, 3p5--> so Cl-= 1s2, 2s2, 2p6, 3s2, 3p6.

Question 38

B) In what order are sub-shells filled up? C) What is the sub-shell notation (electron configuration) of Calcium? D) What is the exception in the order of the sub-shells that get filled up?

Answer 38

B)-1s2 ,2s2,, 2p6, 3s2, 3p6, 4s2, 3d10, 4s2, 4p6, 4d10, 4f14 etc. C)-1s2, 2s2, 2p6, 3s2, 3p6, 4s2. D)-4s sub-shell gas lower energy level than the 3d sub-shell even though it's principal quantum number bigger -->means the 4s sub-shell fills up first.

Question 39

E) Briefly explain how short-hand electron configuration is used

Answer 39

E)-noble gas symbols like that of Argon (Ar), are sometimes used in electron configurations -->E.G: Calcium: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2)-->can be written as--> (Ar) 4s, where (Ar)= 1s2, 2s2, 2p6, 3s2, 3p6.

Question 40

A) Explain why Chromium (Cr) and Copper (Cu) are badly behaved transition metals because they donate one of their 4s electrons to the 3d sub-shell

Answer 40

A)-they donate one of their 4s electrons to the 3d sub-shell - ->this as they are happier with a more stable full OR half-full d sub-shell - ->Cr atom (24 e-): 1s2, 2s2, 2p6, 3s2, 3p6, 3d4, 4s1 - ->Cu atom (29 e-): 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s1.

Question 41

B) Also, what do transitions metals do when they become ions?

Answer 41

B)-they lose their 4s electrons before their 3d electrons | -->Fe atom (26 e-): 1s2, 2s2, 2p6, 3s2, 3p6, 4s2--> Fe 3+ ion (23 e-): 1s2, 2s2, 2p6, 3s2, 3p6, 3d5.

Question 42

A) What determines the chemical properties of an element? B) Describe the ability of the S block elements to form the inert gas configuration

Answer 42

A)-n. of outer electrons of the element B)-these are the ones in groups 1 and 2-->they have 1 or 2 electrons in the outer shell -->these easily lost to form (+) ions with inert gas configuration -->E.G: Na: 1s2, 2s2, 2p6, 3s1-->Na+: 1s2, 2s2, 2p6. (electronic configuration of neon)

Question 43

C) Outline the ability of the P block elements to form the inert gas configuration

Answer 43

C)-elements in group 5, 6 and 7 (p block) can gain 1, 2 or 3 electrons to form (-) ions with inert gas configuration - ->E.G: O--> 1s2, 2s2, 2p4-->O2-: 1s2, 2s2, 2p6 - ->group 4 to 7 can also share electrons when they form covalent bonds

Question 44

D) Describe the reactivity of the GROUP 0 (inert gases) E) Outline the formation of ions from d block elements

Answer 44

D)-have completely filled S and P sub-shells -->so don't need to gain or lose or share electrons -->their full sub-shells make them inert E)-these d block elements are transition metals which tend to lose S and D electrons to form (+) ions.

Question 45

A) What does "ionised" mean? B) Define the term "first ionisation energy"

Answer 45

A)-when electrons have been removed from atom/molecule B)-its the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions -->energy needed to remove first electron.

Question 46

C) Why is ionisation an endothermic process? D) EQUATION: Outline the equation for the first ionisation of oxygen and it's value

Answer 46

C)-have to put energy in to ionise an atom/molecule D)O(g)-->O+(g) + e- -->1st ionisation energy= +1314 KJ/mol-1

Question 47

D) What are the 3 main key points about ionisation energies?

Answer 47

D)1-must always use gas state symbol (g) as ionisation energies measured for gaseous atoms 2-always refer to 1 mole of atoms as stated in definition rather than to single atom 3-lower ionisation energy is-->easier it is for ion to form.

Question 48

A) Explain the following factors and how they affect ionisation energy: 1- Nuclear charge 2- Distance from nucleus 3- Shielding

Answer 48

A)1-more protons in nucleus-->more + charged nucleus is and stronger attraction for electrons 2-attraction falls off v. rapidly with distance -->an electron close to nucleus will be much more strongly attracted than one further away 3-as n. of electrons between outer electrons + nucleus increase the outer electron feel less attraction towards nuclear charge -->this lessening of pull of nucleus by inner shells of electrons called shielding (or screening).

Question 49

B) What does a high ionisation energy mean?

Answer 49

B)-means higher attraction between electron and nucleus | --> and so more energy needed to remove the electron.

Question 50

A) Outline what it means by "successive ionisation energy" B) Define the term "second ionisation energy" C) EXAMPLE: What is the equation for the second ionisation energy of oxygen and the value?

Answer 50

A)-this is energy put in each time an electron is removed -->eventually can remove all electrons to just leave nucleus B)-energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole pf gaseous 2+ ions. c)-EQUATION: O+(g)-->O2+(g) + e- -2nd ionisation energy= +3388 KJ/mol-1

Question 51

D) What is the equation for the "nth" ionisation energy?

Answer 51

D)-X(n-1) (g)--> X (n+) (g) + e-

Question 52

With reference to the graph of successive ionisation energies answer the following: A) How does successive ionisation energies provide evidence of shell structure within atoms?

Answer 52

A)-within each shell successive ionisation energies increase -->as electrons being removed from increasing (+) ion--> less repulsion amongst remaining electrons so they held more strongly by nucleus -big jumps in ionisation energy occur when new shell broken in to--> an electron is being removed from shell closer to nucleus.

Question 53

B) Outline how you would determine which group of the periodic table an element belongs to C) EXAMPLE: What group of the periodic table does Na belong to according to the grpah?

Answer 53

B)-just count n. of electrons removed before first big jump to find the group number c)-1 electron removed before first big jump so Na belongs to group 1.

Question 54

D) Briefly explain how this graph may be used to predict the electronic structure of elements E) EXAMPLE: What is the electronic structure of Na?

Answer 54

D)-working from left to right-->count n.of points there are before each big jump to find n. of electrons in each shell starting with the first E)-graph for Na has 2 points on right-hand side -->then 8 point jump -->and final 1 point -->therefore, Na has 2 electrons in the first shell, 8 in the second and 1 in the third.

Question 55

A) What is the general trend in the first ionisation energies of elements down a group and across a period?

Answer 55

A)-down a group it decreases | -across a period it generally increases

Question 56

Refer to the graph provided: A) Explain why the first ionisation energies decrease down GROUP 2

Answer 56

A)-if each element down group 2 has extra inner electron shell compared to one above-->the extra inner shells will shield outer electrons from attraction to nucleus - also extra shells means that outer electrons are further from nucleus so nucleus's attraction greatly reduced - ->makes sense as both of these factors will make it easier to remove electrons-->results in lower ionisation energies.

Question 57

B) Therefore, what does this trend in ionisation energies down group 2 provide evidence for?

Answer 57

B)-that electron shells do exist and that successive elements down the group have extra, bigger, shells.

Question 58

A) In detail explain why the trend in ionisation energies increases across PERIOD 3 (i.e: harder to remove outer electrons)

Answer 58

A)-this is as n. of protons is increasing-->means stronger attraction to nucleus - ->all extra electrons are at roughly same energy level even if outer electrons are in different orbital types - ->means there's generally little extra shielding effect OR extra distance to lessen attraction from nucleus.

Question 59

Drops in ionisation energies between Groups 2 and 3: EXAMPLE: Magnesium and Aluminium A) What is the electronic figuration of Al and Mg and their 1st ionisation energy values?

Answer 59

A)-Mg: 1s2, 2s2, 2p6, 3s2--> 1st ionisation= 738kj/mol-1 | -Al: 1s2, 2s2, 2p6, 3s2, 3p1--> 1st ionisation 578kj/mol-1

Question 60

B) Explain this drop between group 2 and 3 using the example of Aluminium and Magnesium and what this pattern provides evidence of

Answer 60

B)-Al's outer electron is in 3p orbital rather than 3s - ->3p orbital has slightly higher energy than 3s orbital so electron on average to be found further from nucleus - 3p orbital has additional shielding provided by 3s2 electrons - ->both these factors together strong enough to override effect of increased nuclear charge - ->results in ionisation energy dropping slightly - ->provides evidence for theory of electron sub-shell structure.

Question 61

Drops in ionisation energy between groups 5 and 6: EXAMPLE: Phosphorus (P) and Sulfur (S) A) State the electronic structure and 1st ionisation energy of P and S

Answer 61

A)-P: 1s2, 2s2, 2p6, 3s2, 3p3-->1st ionisation energy= 1012kj/mol-1 -S: 1s2, 2s2, 2p6, 3s2, 3p4-->1st ionisation energy= 1000kj/mol-1

Question 62

B) Explain the drop in ionisation energy between group 5 and 6 C) Once again what does this drop in ionisation energy provide evidence of?

Answer 62

B)1-shielding is identical in phosphorus + sulfur atoms and electron is being removed from identical orbital - in phosphorus's case electron is being removed from singly-occupied orbital - -BUT in sulfur electron is being removed from orbital containing 2 electrons 2-phosphorus: (Ne), 3s2, 3p3 (all 3 singly occupied) -sulfur: (Ne), 3s2, 3p4 (2 singly occupied) -->repulsion between the 2 electrons in an orbital means electrons easier to remove from shared orbitals C)-more evidence for electronic structure model.