Kings College Understanding the Sub Atomic Particles Present in Atoms HW Please check through the work for errors and suggest areas for improvement. I need

Kings College Understanding the Sub Atomic Particles Present in Atoms HW Please check through the work for errors and suggest areas for improvement. I need to get a distinction in this work and so just want to make it perfect before handing it in. Atomic Structure Assignment.
1. Complete the following table.
Name
Symbol
Mass
Charge
Location
Proton
P+
1
+1
nucleus
Neutron
N0
1
0
nucleus
Electron
e-
1/1836
-1
orbitals
2. This question is aimed at understanding the sub-atomic particles present in atoms.
Given the following
14
7
N
16
8
O
32
16
S
39
19
27
13 Al
K
complete the table below.
Particle
Protons
Neutrons
Electrons
Symbol
Sulphur atom
16
16
16
S
Potassium atom
19
20
19
K
Oxygen atom
8
8
8
O
Aluminium ion
13
14
10
AL+3
Nitride ion
7
8
10
N3-
Oxide ion
8
8
10
02-
3. Complete the paragraph regarding isotopes by inserting keywords from the list below into the
shaded boxes.
physical
relative atomic mass
radioactive
mass
element
chemically
nuclei
neutrons
Isotopes are atoms of the same
element
With different numbers of
neutrons
They react chemically
In exactly the same way, but their
physical
properties are different due to atoms of different isotopes having different
mass
Sometimes the
nuclei
of isotopes are unstable, which results in the element
becoming
radioactive
Nearly all elements exist in isotopic form, but often there
Is one main isotope. Accordingly, the
Relative atomic mass
is very close to the mass
number of that particular element.
4. The table below contains information about several isotopes. Use the information given to fill in
the blanks. Assume all atoms are neutral. (note these are isotopes!!)
Isotope
Nuclear
Atomic
Mass
No. of
No. of
No. of
Name
Symbol
Number
Number
Protons
Neutrons Electrons
40
calcium
55Manganese
Oxygen 18
Gold 197
40CA
20
40
20
20
20
55MN
25
55
25
30
25
18O
8
18
8
10
8
79
197
79
118
79
197
79
Au
5. The table below contains information about several ions. Use the information given to fill in the
blanks.
Element
Ion
Atomic
Mass
No. of
No. of
No. of
Name
Symbol
Number
Number
Protons
Neutrons Electrons
Fluoride
F?
9
19
9
10
10
Cadmium
CD+2
48
112
48
64
46
Oxide
02-
8
16
8
8
10
Aluminium
Al 3+
13
27
13
14
10
6. Calculate the relative atomic mass for Chlorine if a sample, passed through a mass spectrometer,
yields one peak with a mass value of 35a.m.u and a relative abundance of 75% and a second peak
with a mass value of 37 a.m.u and a relative abundance of 25%.
(35×75)+(37×25)/100=35.5
7. Calculate the relative atomic mass for neon if its abundance in nature is 90.5% 20 neon, 0.3%
21
neon, and 9.2% 22neon.
(90.5×20)+(0.3×21)+(9.2×22)/100=20.187
8. Calculate the relative atomic mass of silver if 13 out of 25 atoms are 107silver and 12 out of 25
atoms are 109silver.
(13×107)+(12×109)/100=26.99
9.
The mass of an atom is contained mainly in its
nucleus
10. The identity of an element is determined by its number of
11. Isotopes are atoms with the same number of
protons
12. The charge of an atom or ion is determined by its number of
and
electrons
protons
and different number of
neutrons
electrons
13. Particle X contains 9 protons, 10 neutrons, and 9 electrons. Particle Y contains 9 protons, 10
neutrons, and 10 electrons. What is the relationship between particles X and Y? In the shaded boxes
below mark ‘T’ for true and ‘F’ for false.
A.
B.
C.
D.
Particles X and Y are isotopes of the same element.
Particle X is an atom, and particle Y is an ion of the same element.
Particle X and Y are atoms of different elements.
There is no significant difference between particles X and Y.
T
T
F
F
14. Briefly explain your answers to question 13 in the box below.
A. They are both isotopes of Flourine . B. By adding 1 electron it becomes an electrically charged particle
(negative ion) .C. They aren’t atoms of different elements because Flourine is the only element with 9
protons. D Because of the above answers there are clear differences between X and Y.
15. In the box below, describe the shapes of the s, p and d orbitals.
1. The S Orbital is Spherical and symmetrical around the nucleus of the atom , The 2s orbital is very
similar to the 1s orbital except the area where there is the greatest chance of finding an electron is
further from the nucleus, 3s orbitals are also spherical and have gaps in them called nodes. 2. Unlike
an S orbital P orbitals point in different directions and look like a balloon which has been tied at the
centre. P orbitals are shaped like a pair of lobes on opposite sides of the nucleus. .3. D orbitals are
shaped like clover leaves.
16. Label the following orbitals;
S
py
pz
px
17. Compare and contrast Hund’s rule with the Pauli exclusion principle.
Both theories describe the electrons and orbitals in atoms. Hunds rule put simply is that every Orbital in a
subshell is singly occupied with one electron before any one orbital is doubly occupied, all electrons in
singly occupied orbitals have the same spin. Pauli’s exclusion principle states that no two electrons in the
same atom can have identical values for all four of their quantum numbers, so no more then 2 electrons
can occupy the same orbital, two electrons in the same orbital must have opposite spins. PEP is about
quantum numbers of an atom. Hund rule is about how electrons are filled to the orbitals of an atom. PEP
says of having only 2 electrons per orbitals and Hund says that only after filling one electron to each
orbital the electron pairing starts to happen. PEP describes how electrons in the same orbitals have
opposite spins, which can be used to explain the Hund rule.
18. How does the figure below illustrate Hund’s rule?
??
2p
2s
??
1s
??
?
?
Each electron in singly occupied orbitals have the same spin and orbitals of the same energy are filled
first
19. How does the figure above illustrate the Pauli Exclusion Principle?
Each orbital containing pairs of electrons shows that the electrons have opposite spin.
20. Electron Configurations are written in the following form. Describe what each part represents.
Quantum number
Electrons
1s2
Orbital (s sub shell)
21. State (yes/no) whether each of the following orbital diagrams conforms to the rules governing
electron configuration. If not, explain what is wrong with the diagram. Note: only valence
electrons are shown.
??
2s
a.
??
b.
2p
??
2s
?
?
2p
?
a.
No, Each orbital would be filled before any of them has 2 electrons
c.
??
2s
?
?
2p
?
b.
enter text.
No, The electrons in the 2s orbital would have opposite spins
c.
enter text.
Yes, Each orbital has at least one electron before the 2s orbital gained another and the 2s orbital
contains 2 electrons with opposite spins which is correct.
22. Identify the following atoms by their orbital diagrams and write the electron configuration
??
_____
1s
??
_____
2s
Atom: P
??
_____
1s
??
_____
1s
Electron configuration
??
_____
2s
Atom: Na
?
?
?
_____ _____ _____
3px
3py 3pz
?
_____
3s
[Ne]3s1
??
??
??
_____ _____ _____
2px
2py
2pz
Electron configuration
??
_____
3s
[Ne]3S23P3
??
??
??
_____ _____ _____
2px
2py
2pz
Electron configuration
??
_____
2s
Atom: K
??
??
??
_____ _____ _____
2px
2py
2pz
??
_____
3s
??
??
??
_____ _____ _____
3px
3py 3pz
[Ar]4s1
23. Identify the missing orbitals. Place your answers in the shaded boxes.
5p
?
_____
4s
4d
5s
4p
3d
4s
3p
3s
2p
2s
1s
24. Write out the full electron arrangement for Sodium.
1s22s22p63s1
or
[Ne]3s1
25. Write the electron arrangement of Bromine, using an inert gas symbol as shorthand.
enter text.
[Ar]3d104s24p5
26. Write out the electron configuration for a Ca2+ ion using an inert gas symbol as shorthand.
[Kr]4d10
Ionisation energies
This page is to test your understanding of ionisation energies. The graph below shows the first
and second ionisation energies of the elements neon to calcium.
27. Why is the second ionisation energy for each element greater than its corresponding first
ionisation energy?
An elements second ionization energy is the energy needed to remove the outermost (least bound)
electron from a 1+ion of the element. Because positive charge binds electrons more strongly the
second ionisation energy is always higher then the first.
28. Explain why Group I elements have the lowest first ionisation energy but the highest second
ionisation energy?
When you remove an electron, you need a certain amount of energy, or the first ionisation energy. This is
a certain amount of energy depending on the element, and when you have taken the first electron, you
leave a positive ion behind, which makes a stronger attraction on the remaining electrons.Group one
elements (the Alkali metals) can be referred to as s-block elements because their highest energy
electrons appear in the s subshell. Going down the group the first ionisation energy decreases. There is
more shielding between the nucleus and the outer electrons and the distance between the nucleus and
the outer electron increases and therefore the force of attraction between the nucleus and outer most
electrons is reduced meaning that the first ionisation energy is low as an electron is lost easily.
However once the electron is lost the group 1 elements receive a +1 charge after the first ionization
and the positive charge binds the electrons more strongly meaning that the second ionisation energy
will be higher. In the group 1 elements there is more shielding between the nucleus and the outer
electrons and the highest energy electrons appear in the s-subshell.
29. Explain why neon and argon have the highest first ionisation energies?
They belong to group 18 , the noble gases and have a very stable electron configuration. They have full
outer shell of valance electrons making the removal of an electron extremely difficult. They also have
the most protons of their period so atomic radius is reduced as electrons are pulled into the nucleus.
30. Why is the first ionisation energy for aluminium less than for magnesium?
Magnesium has its outermost electrons in the 3s sub-level. The aluminium atom has its outermost
electron in the 3p sublevel. Since p electrons have just slightly more energy than s electrons, it takes a
little less energy to remove that electron from aluminium. One other factor is that the electrons in 3s2
shield the electron in 3p1. These two factors allow the IE1 for aluminium to be less than IE1 for
magnesium.
31. Why is the first or second ionisation energy for calcium lower than the first or second ionisation
energy for magnesium?
The magnesium S orbital is fully filled and therefore more stable.
A
B
C
D
E
F
G
H
Sodium
Magnesium
Aluminum
Silicon
Phosphorus
Sulfur
Chlorine
Argon
32. For the graph of first ionisation energies below;
(a) Label the points A-H with the correct element;
(b) explain the down turn at B-C.
The 3p electron of aluminium is further from the nucleus compared to the 3s electrons of magnesium.
SO…27,28,29,30,31+32 all need to be way more detailed

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