Atomic theory 1 Flashcards
1
Q
Naturally occurring chlorine consists of 75.5% of (35)Cl and 24.5% of (37) Cl. Calculate the relative atomic mass (Ar) of chlorine. (L.C)
A
- 75.5 x 35 = 2642.5
- 24.5 x 37 = 906.5
100 atoms = 3549
1 atom = 35.49 (Ar)
2
Q
A dipositive ion, M2+, has 25 electrons and 32 neutrons. What is (i) the atomic number, (ii) the mass number, of M (L.C)
A
i) 27
ii) 59
3
Q
What are isotopes? (L.C)
A
Atoms with the same atomic number (Z) but different mass numbers (A)
4
Q
How many (i) electrons, (ii) neutrons, has the aluminium ion (13 - bottom, 27 - top) Al³+ (L.C)
A
i) 10
ii) 14
5
Q
Explain why relative atomic masses are rarely whole numbers. (L.C)
A
As it is the average of mass numbers of the isotopes of an element
6
Q
Define (i) mass number, (ii) relative atomic mass (L.C)
A
(i) number of protons and neutrons in the atoms of an isotope
(ii) Average mass of atoms of element relative to 1/12 of mass of carbon-12 atom
7
Q
A sample of the element gallium is composed of 60.1% gallium-69 and 39.9% gallium-71. Calculate the relative atomic mass of gallium from this information. (L.C)
A
69 x 60.1 = 4146.9
71 x 39.9 = 2832.9
100 atoms = 6979.8
Ar = 69.798 (69.8)
8
Q
Define relative atomic mass (L.C)
A
Average mass of atoms of element relative to 1/12 mass of carbon-12 atom
9
Q
Define relative atomic mass (L.C)
A
Average mass of atoms of element relative to 1/12 mass of carbon-12 atom
10
Q
Define (a) atomic number, (b) relative atomic mass (L.C)
A
a) Number of protons in the nucleus of an atom of the element
b) Average mass of atoms of element relative to 1/12 mass of carbon-12 atom
11
Q
What are isotopes? (L.C)
A
Atoms with the same atomic number (Z) but different mass numbers (A)
12
Q
What are isotopes? (L.C)
A
Atoms with the same atomic number (Z) but different mass numbers (A)
13
Q
What is the principle of the mass spectrometer? (L.C)
A
Positive ions separated based on relative mass when moving in a magnetic field
14
Q
Calculate, to two decimal places, the relative atomic mass of a sample of neon shown by mass spectrometer to be composed of 90.5% of neon-20 and 9.5% of neon-22. (L.C)
A
90.5 x 20 = 1810
9.5 x 22 = 209
100 atoms = 2019
Ar = 20.19
15
Q
What are isotopes? (L.C)
A
Atoms with the same atomic number (Z) but different mass numbers (A)
16
Q
Define relative atomic mass, Ar (L.C)
A
Average mass of atoms of element relative to 1/12 mass of carbon-12 atom
17
Q
Calculate the relative atomic mass of a sample of lithium, given that a mass spectrometer shows it consists of 7.4% (6)Li and 92.6% (7)Li (L.C)
A
7.4 x 6 + 92.6 x 7 = 692.6
100 atoms = 692.6
1 atom = 6.926
18
Q
What is the principle on which the mass spectrometer in based? (L.C)
A
Positive ions separated based on relative masses when moving in a magnetic field.
19
Q
Identify an element that is a non-metal and is a liquid at room temperature.
A
Br
20
Q
Identify an element that is a divalent metal.
A
Be (anyone from group 2)
21
Q
In the periodic table identify an element in the same period as magnesium but with larger atoms. (L.C)
A
Sodium / Na
22
Q
In the periodic table identify an element in the same period as magnesium but with smaller atoms. (L.C)
A
Beryllium / Be
23
Q
Define electronegativity
A
The relative power of attraction of an atom of an element for the shared pair of electrons in a covalent bond.
24
Q
Describe using dot and cross diagrams the bonding in the ammonia molecule.
A
Drawn.
25
Use electron pair repulsion theory to determine the shape of the ammonia molecule. Which of the following angles: 104°, 107°, 109°, 120° or 180° would you expect to be closest to the bond angle in the ammonia molecule? Explain your answer.
- Three bonding and one non-bonding (lone) pair giving bond arrangement (shape of molecule) to be pyramidal
- 107°
- Greater repulsion of lone pair and lone pair pushes bonds closer together
26
Define electronegativity.
The relative power of attraction of an atom of an element for the shared pair of electrons in a covalent bond.
27
Account for the increase in electronegativity values across the second period of the periodic table.
- Effective nuclear charge increasing
- More protons
- Atomic number (Z) increasing /
- Atomic radius decreasing
28
Use electronegativity values to predict the type of bonding in oxygen difluoride (OF₂).
Slightly polar covalent
29
State and account for the shape of the OF₂ molecule.
State - V-shaped planar
| Account - 2 bond pairs electrons and 2 lone pairs
30
180.0° 109.5 ° 120.0° 103.0° Select, giving your reasons, which of these angles is the most probable value for the bond angle in oxygen difluoride.
xx
31
What is meant by the periodic system in the context of Mendeleev’s 1869 periodic table of the elements?
Elements listed according to relative atomic mass and in groups with similar properties
32
Comment on the positioning of tellurium (Te) and iodine (I) in the 1869 table.
Chemical properties matched (fitted) better when order reversed.
33
Why did the 1869 table not include any noble gases?
They were undiscovered in 1869.
34
Modern periodic tables arrange the elements in order of atomic number and not in order of relative atomic mass. Define the terms atomic number and relative atomic mass.
A.N - The number of protons present in the nucleus of an atom, (this equals the number of electrons in an atom).
R.A.M - The average mass of an atom of an element compared to 1/12 the mass of the carbon - 12 isotope, taking relative abundances of isotopes into account.
35
Explain why all the elements of Group 18 in the periodic table are chemically inert.
Stable arrangement of electrons.
36
Explain how and why the reactivity of the halogens
| changes down Group 17.
How : Less reactive down group
| Why : Increasing (greater) atomic radius (number of shells)
37
How would you expect a small sample of francium to react in water? Justify your answer. Predict the products of this reaction.
How : violently / explosively
Justify : reactivity increases down group
Predict : FrOH (francium hydroxide) and H₂ (hydrogen)
38
State and give the reason for the trend in atomic radii across the second period of the periodic table.
State : decrease
| Give : nuclear charge increasing and atomic number increasing
39
Define (i) mass number of an atom, (ii) relative atomic mass of an element.
(i) The number of protons plus neutrons in the nucleus of an atom.
(ii) The average mass of an atom of an element compared to 1/12 the mass of the carbon - 12 isotope, taking relative abundances of isotopes into account.
40
A sample of magnesium metal was introduced into a mass spectrometer and vaporised. What were the next three fundamental processes that occurred in the spectrometer?
- ionisation
- acceleration
- separation
41
Define electronegativity.
The relative power of attraction of an atom of an element for the shared pair of electrons in a covalent bond.
42
Predict the type of bond formed between carbon and chlorine atoms in a CCl₄ molecule.
Polar covalent
43
What is the valency of carbon in tetrachloromethane?
4
44
State and account for the shape of a tetrachloromethane molecule.
State : tetrahedral
| Account : four pairs electrons and no lone pairs.
45
Draw a dot and cross diagram to show the arrangement of all the valence shell electrons in a CS₂ molecule.
Drawn.
46
Why did Mendeleev place tellurium before iodine in his periodic table of elements?
x
47
Define electronegativity.
The relative power of attraction of an atom of an element for the shared pair of electrons in a covalent bond.
48
Why is there an increase in electronegativity value moving from gallium to germanium in the periodic table?
- nuclear charge increasing
| - atomic radius decreasing
49
Mendeleev predicted the properties of the elements gallium and germanium years before either of them was discovered. Explain the basis for his predictions.
Predicted the properties from properties of known elements.
50
Write the molecular formula for the simplest compound formed between germanium(IV) and hydrogen. Would you expect this compound to be water soluble? Justify your answer.
GeH₄
Would : no
Justify : GeH₄ a non-polar (slightly polar) solute
51
In the periodic table, identify an element in the same period as magnesium but with larger atoms.
Sodium (Na)
52
In the periodic table, identify an element in the same group as magnesium but with smaller atoms.
Beryllium (Be)
53
Define first ionisation energy.
This is the minimum amount of energy required to remove the first most loosely bound electron from a mole of isolated atoms of an element in its neutral gaseous ground state. X(g) → X⁺(g) + e-
54
Explain why the first ionisation energy value of silicon is (i) greater than that of aluminium, (ii) less than that of carbon.
(i) greater nuclear charge and smaller atomic radius
| (ii) greater atomic radius
55
What experimental evidence do we have other than graphs for the existence of energy levels in atoms?
Line emission (absorption) spectra of elements
56
Define first ionisation energy.
This is the minimum amount of energy required to remove the first most loosely bound electron from a mole of isolated atoms of an element in its neutral gaseous ground state. X(g) → X⁺(g) + e-
57
Define bond energy.
The average energy required to break a bond (to break 1 mole of bonds) and to separate the atoms into separate single atoms or in the gaseous state.
58
Define atomic radius (covalent radius).
half internuclear distance (half distance between the centres of the atoms) in a single homonuclear bond (of singly-bonded atoms of the same element)
59
State and explain the trend in atomic radii (covalent radii) across the second period of the periodic table of the elements.
State : decrease in atomic radius
| Explain : increase in effective nuclear charge (number of protons)
60
Give two reasons why electronegativity values exhibit a general increase across the second period of the periodic table.
- increase in nuclear charge (number of protons)
| - decrease in atomic radius
61
H₂O, NH₃, PH₃, HCl
State how the bonding in PH3 differs from the bonding in the other three hydrides. What is the reason for this difference in bonding?
State : PH₃ virtually non-polar (pure covalent) but the other three are polar covalent.
What: Tiny electronegativity difference in PH₃ (between P and H), but a much bigger electronegativity differences in the other three.
62
H₂O, NH₃, PH₃, HCl
From these four hydrides, identify the hydride or hydrides in which hydrogen bonding occurs between the molecules. Give one property that is affected by the presence of intermolecular hydrogen bonding in the hydride or hydrides you have identified.
- H₂O and NH₃
| Give - melting point / boiling point
63
State the shape of the PH₃ molecule and explain using electron-pair repulsion theory how this shape arises.
Shape : pyramidal
| Explain : repulsion between four electron pairs (e.p.), one a lone pair (l.p.)
64
Boron trichloride (BCl₃) is a colourless gas. Would you expect (i) the B–Cl bonds, (ii) the BCl₃ molecules, to be polar or non-polar? Justify your answers.
(i) B – Cl bond: polar
(ii) BCl3 molecule: non-polar
unequal sharing of electrons (el. neg. difference) between B and Cl (polarity of bonds) cancels due to symmetry of molecule(s)
65
Define (i) atomic number, (ii) relative atomic mass.
(i) The number of protons present in the nucleus of an atom, this equals the number of electrons in an atom.
(ii) The average mass of an atom of an element compared to 1/12 the mass of the carbon -12 isotope, taking relative abundances of isotopes into account.
66
What was the basis (periodic law) used by Mendeleev in arranging the elements in his periodic table?
When arranged in order of increasing atomic weight (relative atomic mass) there is a periodic occurrence of (elements with similar properties.
67
Why did Mendeleev leave spaces in his periodic table, e.g. where the element germanium occurs in the modern periodic table?
So that similar elements (elements with same properties) were in same group
68
In a few instances Mendeleev reversed the order of elements required by his periodic law. For example, he placed the element tellurium before the element iodine. Explain why he did this.
To suit (fit) properties to groups
69
Explain why (i) the alkali metals are all reactive, (ii) the reactivity of the alkali metals increases down the group.
(i) Readily lose an electron in outer shell and have a low first ionisation energy
(ii) An increase in atomic radius
70
Define atomic orbital.
An atomic orbital is a region in space around the nucleus of an atom in which there is a high probability of finding an electron.
71
Write the electron configuration (s, p, etc.) of the element manganese (Mn).
1s₂2s₂2p₆3s₂3p₆4s₂3d₅
72
What do the electron configurations of the series of elements from scandium to zinc have in common?
- Electrons entering (occupying) 3d sublevel
| - All end in 3d(x)
73
Define electronegativity.
The relative power of attraction of an atom of an element for the shared pair of electrons in a covalent bond.
74
State two factors that cause electronegativity values to increase across a period in the periodic table of the elements.
- Increasing effective nuclear charge
| - Decreasing atomic radius
75
State two differences between Mendeleev's periodic table and the modern periodic table of the elements.
Arranged in terms of atomic weight, reversed some pairs of elements and left gaps for undiscovered elements.
76
Account for the trend in the size of the atomic radius going across a period in the periodic table.
Across - decrease in atomic radius
| Account - increase in nuclear charge
77
Write the electron configuration (s, p, etc.) of the aluminium ion (Al³⁺).
1s₂2s₂2p₆
78
What contribution did Henry Moseley make to the systematic arrangement of the elements in the periodic table?
Found the number of protons in nucleus
79
Give two properties of alpha particles.
Positive (+) charge (attracted to negative) and poor penetration.
80
Define electronegativity.
The relative power of attraction of an atom of an element for the shared pair of electrons in a covalent bond.
81
State and explain the trend in electronegativity values down the first group in the periodic table of the elements.
State - Decrease
| Explain - increasing atomic radius and effective nuclear charge unchanged (constant)
82
Use electronegativity values to predict the types of bonding (i) in water, (ii) in methane, (iii) in magnesium chloride.
(i) water - polar covalent
(ii) methane - covalent
(iii) magnesium chloride - ionic (electrovalent)
83
Define atomic (covalent) radius.
Half the distance between the centres of singly-bonded atoms of the same element.
84
Use electronegativity values (log tables) to predict the type of bond expected between hydrogen and sulfur.
- Weakly polar
- Almost non-polar
- Covalent bond
85
What contribution did Newlands make to the systematic arrangement of the elements known to him?
Arranged in increasing relative atomic mass (atomic weight) and in his law of octaves.
86
Define electronegativity.
The relative power of attraction of an atom of an element for the shared pair of electrons in a covalent bond.
87
Explain why there is a general increase in electronegativity values across the periods in the periodic table of the elements.
- Decrease in atomic radius
| - Increase in effective nuclear charge
88
Explain, in terms of the structures of the atoms, the trend in reactivity down Group 1 (the alkali metal group) of the periodic table.
- Reactivity increases
- Increase in atomic radius
- Effective nuclear charge is the same
89
The minimum energy required to completely remove the most loosely bound electron from a mole of gaseous atoms in their ground state defines an important property of every element. Identify the energy quantity defined above. State the unit used to measure this quantity.
Identify : first ionisation energy
| State : kilojoules per mole (kJ mol–1)
90
Using X to represent an element, express the definition first ionisation energy in the form of a balanced chemical equation.
X(g) → X⁺(g) + e–
91
Would it take more or less energy to remove the most loosely bound electron from an atom if that electron were not in its ground state? Explain.
Would : less
| Explain : already gained energy / already excited
92
An element has a low first ionisation energy value and a low electronegativity value. What does this information tell you about how reactive the element is likely to be, and what is likely to happen to the atoms of the element when they react?
How : reactive
| React : by losing electron(s)
93
Define electronegativity.
The relative power of attraction of an atom of an element for the shared pair of electrons in a covalent bond.
94
What contribution did Dobereiner make to the systematic arrangement of the elements?
He grouped elements of similar properties in groups of three (triads).
95
Define atomic radius (covalent radius).
Half the distance between the centres of singly-bonded atoms of the same element.
96
Describe and account for the trend in atomic radii (covalent radii) of the elements (i) across the second period, (ii) down any group, of the periodic table.
(i) Across : decrease in atomic radius
Account : increase in nuclear charge and increase in effective nuclear charge
(ii) Down : increase in atomic radius
Account : increase in number of filled shells
97
(In the graph 2003) Account for the general increase in ionisation energies across these elements.
- Increase in nuclear charde
- Increase in atomic number
- Decrease in atomic radius
98
(In the graph 2003) Explain why the ionisation energies of element number 4 and 7 are exceptionally high relative to the general trend.
Element 4 : The outer sublevel is full giving increased stability.
Element 7 : A half-full p sublevel giving increased stability.
99
How does the definition of second ionisation energy differ from that of first ionisation energy.
Electron removed from monopositive ion.
100
State two ways in which Mendeleev's periodic table of the elements differs from that of Moseley.
Mendeleev : In order of atomic mass
Moseley : In order of atomic number
Mendeleev : Reversed elements
Mosely : No reversing
101
Define first ionisation energy.
The minimum energy to remove most loosely-bound electron from an isolated atom in its ground state.
102
Account fully for the trends in first ionisation energies of elements across the second period of the periodic table (i.e. Li to Ne).
```
Increase:
- increasing nuclear charge (atomic number)
- decreasing atomic radius
Drops :
- Be or N stable
- Be due to outer sublevel full
- N due to half-full p sublevel
```
103
Account for the trend in first ionisation energies of the elements going down Group II of the periodic table, i.e. the alkaline-earth metals.
- decrease due to increasing atomic radius,
- increased shielding (screening) offsets
- increased nuclear charge
- effective nuclear charge remains constant
104
(From graph 2002 q.5) Explain why there is an increase in these ionisation energy values.
- Increasing positive charge of species losing electron (no. of protons same but no. of electrons decreasing) - Decreasing radius
105
(From graph 2002 q.5) Account for the dramatic increase in ionisation energy going from the second to the third ionisation. Between which two ionisations would you expect the next dramatic increase to occur if the data for further ionisation energies of magnesium were examined? Give a reason for your answer.
2nd – 3rd: third electron coming from new shell (new main level)
The next: between 10th and 11th
Reason : eleventh electron coming from new shell (new main level)
106
The bond energy values in kJ mol-1 of the C-CI bond, the C-Br bond and the C-I bond are 338, 276 and 238 respectively. Suggest a reason for this trend.
x
107
(In graph 2001 Q.4) Which of the elements shown, shows behaviour closest to that of an ideal gas?
Helium
108
Explain why atomic radius decreases going from aluminium to chlorine.
- Increasing nuclear charge
109
Explain why atomic radius increases going from nitrogen to phosphorus.
- Idea of extra shell of electrons
110
Explain why the atomic radii of chromium, copper and zinc are very similar?
- No change in effective nuclear charge
111
One of the elements, whose symbol is given in the table above, melts at 172 K. At temperatures above 239 K it is a gas which is denser than air. Identify the element and state a common use for it. Write an equation for the reaction of water with the compound formed between this element and aluminum.
Chlorine / Cl
Use : Water purification
Equation : ***************************************
112
Define electronegativity. Both nitrogen and phosphorus combine with hydrogen to form gaseous hydrides. Which of these hydrides would you expect to be readily soluble in water? Give reasons for your choice.
x
113
Write the electronic configurations (s, p, etc.) for chromium and copper and explain why these two elements are exceptions to the normal order in which electrons occupy sublevels.
x
114
Chromium, copper and zinc are all classified as d-block elements but only chromium and copper are classified as transition elements. Explain why this is the case.
Transition : Cr, Cu can have incomplete d and Zu has complete d.
115
Give an advantage of arranging the elements in order of atomic number.
No need to reverse order to force elements into correct groups.
116
Define relative atomic mass number.
The average mass of an atom of an element compared to 1/12 the mass of the carbon - 12 isotope, taking relative abundances of isotopes into account.
117
Calculate the relative molecular mass of the element X whose composition was found to be 70% 63X and 30% 65 X. Identify element X.
70% x 63 = 44.1
30% x 65 = 19.5
+ = 63.6 = Cu / Copper
118
How many (a) protons and (b) neutrons are present in 12 27Mg 2+
(c) electrons = 10
| (d) neutrons = 15
119
Identify the scientist associated with the discovery of (i) the proton, (ii) the size of the charge of an electron, (iii) triads.
(i)
(ii) Milikan
(iii) Dobereiner
120
How did Rutherford deduce that the nucleus was (i) dense and (ii) small.
x
121
How many (i) electrons and (ii) neutrons n 34S 2-?
x
122
What is the relative atomic mass of magnesium which normally consists of 78.6% magnesium-24, 10.1% magnesium-25 and 11.3% magnesium-26?
x
123
If M3+ has 15 electrons, identify M.
Chlorine / Cl
124
What is an isotope?
Atoms with the same atomic number (Z) but different mass numbers (A).
