Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Chemistry Exam Questions > Topic 1: Atomic Structure and the Periodic Table > Flashcards

Topic 1: Atomic Structure and the Periodic Table Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

Hydrogen, deuterium and tritium are all isotopes of each other.

Identify one similarly and one difference between these isotopes. [2]

A

Similarity:
1. They’ve all got the same number of protons/electrons;

Differences:
2. They all have different numbers of neutrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Hydrogen, deuterium and tritium are all isotopes of each other.
2
Deuterium can be written as 1H. Determine the number of protons, neutrons and electrons. [1]

A
  1. 1 proton, 1 neutron, 1 electron
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Hydrogen, deuterium and tritium are all isotopes of each other.

Write the nuclear symbol for tritium, given that it has 2 neutrons. [1]

A
  1. 3
    H
    1
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Explain why the relative atomic mass of copper (2 isotopes) is not a whole number. [2]

A
  1. A sample of copper is a mixture of isotopes in different abundances;
  2. The relative atomic mass is an average mass of these isotopes which isn’t a whole number
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Give the electron configuration of a Cu atom (29 electrons)

A

1s2 2s2 2p6 3s2 3p6 3d10 4s1

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

What happens in an atom when energy is emitted? [2]

A
  1. Electrons move;
  2. From higher to lower energy levels
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Explain why the lines in a set of lines on an emission spectrum get closer together as you go along. [1]

A
  1. The energy levels (shells) get closer together with increasing energy
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What happens when an electron moves from a higher to a lower quantum shell? [1]

A
  1. Energy is released/emitted
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Describe what the lines on an emission spectrum show. [1]

A
  1. Frequencies of light emitted when an electron drops to a lower energy level
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Explain how emission spectra provide evidence that supports our current understanding of electrons existing in fixed energy levels [2]

A
  1. Emission spectra show that specific amounts of energy are emitted when electrons drop down from higher energy levels to lower energy levels;
  2. In-between amounts of energy are never emitted, which suggests that electrons only exist at very specific energy levels (they’re discrete)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Write an equation, including state symbols, to represent the first ionisation energy of carbon (C). [2]

A

C(g) —) C+(g) + e-

  1. Correct equation;
  2. Correct state symbols
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Explain why increasing nuclear charge increases first ionisation energies. [2]

A
  1. As the nuclear charge increases there is a stronger force of attraction between the nucleus and the electron;
  2. So more energy is needed to remove the electron
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Why does it take more energy to remove each successive electron? [2]

A
  1. The electrons are being removed from an increasingly positive ion;
  2. and there’s less repulsion among the remaining electrons so they’re held more strongly by the nucleus
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What causes the sudden increases in ionisation energy? [1]

A
  1. When an electron is removed from a different shell there is a big increase in the energy required (since that shell is closer to the nucleus)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Explain why the melting point of sulphur is higher than that of phosphorus. [2]

A
  1. Sulphur (S8) has more electrons than phosphorus (P4);
  2. Which results in stronger London forces of attraction between molecules
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What does emission spectroscopy provide evidence for the existence of? [1]

A
  1. Quantum shells
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

What is the electronic configuration for the S2- ion? (S = 16e) [1]

A
  1. 1s2 2s2 2p6 3s2 3p6
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Which is the most likely sequence of values, in kJ mol-1, for the first four ionisation energies of barium?

A 1000 2251 3361 4564
B 496 4563 6913 9544
C 503 965 3458 4530
D 578 1817 2745 11578 [1]

A
  1. C
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Which list contains only s-block elements?

A Li, Na, Mg, Cl
B K, Ca, Co, Rb
C Mg, Al, Sr, Ba
D Be, Rb, Ba, Ra [1]

A
  1. D
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

What is the electronic configuration of the arsenide ion, As3-? (As = 33e) [1]

A
  1. 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

What is the total number of electrons in all the occupied p orbitals in a chloride ion, Cl-? (Cl = 17e) [1]

A
  1. 12
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

How many of each subatomic particle are there in oxygen-18? [1]

A
  1. 8 protons, 10 neutrons, 8 electrons
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

Which of the following sets of data showing the first four ionisation energies, in kJ mol-1, of four elements is most likely to belong to boron?

A 1086, 2353, 4621, 6223
B 900, 1757, 14849, 21007
C 801, 2427, 3660, 25026
D 578, 1817, 2745, 11578 [1]

A
  1. C
24
Q

Which of these ions has the electronic configuration [Ar]3d5?

A Cr3+
B Fe2+
C Mn2+
D Mn3+ [1]

A
  1. C
25
Q

What is the equation for the second ionisation of bromine?

A

Br+(g) —) Br2+(g) + e-

26
Q

Which of the following pairs of ions is isoelectronic?

A N3- and Cl-
B O2- and S2-
C Na+ and K+
D Na+ and Mg2+ [1]

A
  1. D
27
Q

What is the electronic configuration of a bromine atom (35e-)? [1]

A
  1. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
28
Q

What is the electronic configuration of an aluminium atom (13e-)? [1]

A
  1. 1s2 2s2 2p6 3s2 3p1
29
Q

How many electrons are in the fourth quantum shell of bromine (35e-)?

A
  1. 7
30
Q

State what is meant by two arrows in a box when showing electron configuration. [1]

A
  1. (Up and down arrows represent) two electrons in the same orbital with opposite spins
31
Q

Chlorine and iodine are in the same group in the Periodic Table.

Explain why iodine and chlorine have many similarly chemical reactions. [2]

A
  1. Iodine (also) has 7 electrons in the outer shell;
  2. Electronic configurations in the outer shell determines their chemical reactions
32
Q

What does the bonding in magnesium result from? [1]

A
  1. Strong electrostatic attractions between positively charged ions and a sea of delocalised electrons
33
Q

Explain why the first ionisation energy of magnesium is higher than that of sodium. [3]

A
  1. Magnesium has a greater nuclear charge;
  2. Shielding in magnesium atom similar to that in sodium atom;
  3. So the force of attraction between the nucleus and the (outer) electron is greater in magnesium
34
Q

Write the equation, including state symbols, to show the third ionisation energy of magnesium. [1]

A

Mg2+(g) —) Mg3+(g) + e-

35
Q

A student suggested that the difference in the rates of reaction of strontium and barium with water is due to the difference in the sum of their first and second ionisation energies. Discuss this suggestion
[6]

A
  1. Barium loses electrons more easily;
  2. Barium has a larger atomic radius;
  3. Barium has more shielding;
  4. This overrides the influence of barium having (…);
  5. a greater nuclear charge;
  6. Barium reacts faster
36
Q

What is the equation for the first ionisation energy of sulphur? [1]

A
  1. S(g) —) S+(g) + e-
37
Q

Explain the trend in the values of the first ionisation energies for Group 6. [3]

A
  1. First ionisation energy decreases down the group because, although the number of protons is increasing,;
  2. the electron being removed is further from the nucleus;
  3. giving more shielding from the nucleus
38
Q

Explain why the first ionisation energy of sulphur is lower than that of phosphorus. [2]

A
  1. In sulphur electron being removed from an orbital containing two electrons;
  2. (increasing) repulsion between electrons (so the electron is lost more easily)
39
Q

Explain why the first ionisation energy of sulphur is lower than that of chlorine. [2]

A
  1. In sulphur the nuclear charge is lower;
  2. And the electron being removed has similar shielding
40
Q

Explain why the first ionisation energy of hydrogen is less than that of helium, but greater than that of lithium. [4]

A
  1. He has a greater nuclear charge than H;
  2. In He the outer electrons is in the same shell as hydrogen;
  3. In Li the outer electron is further from the nucleus;
  4. (and) is shielded by inner electrons
41
Q

A sodium atom has 11 protons whereas a potassium atom has 19 protons.

Explain why the first ionisation energy of sodium is greater than that of potassium. [3]

A
  1. The outer electron in a sodium atom is closer to the nucleus;
  2. (and) less shielding from inner electron shells;
  3. These outweigh the greater nuclear charge in potassium
42
Q

Explain why the difference between the second and third ionisation energies of calcium is much larger than the difference between the first and second ionisation energies. [2]

A
  1. The 3rd electron is lost from a shell closer to the nucleus;
  2. 1st and 2nd electrons removed from the same shell
43
Q

Give a reason why successive ionisation energies increase. [1]

A
  1. Electron is removed from an increasingly positive charged ion

OR

Electron removed is closer to the nucleus

OR

Electron removed is experiencing less electron-electron repulsion

44
Q

State what is meant by the term first ionisation energy. [3]

A
  1. Energy required to remove an electron;
  2. from 1 mole;
  3. of gaseous atoms
45
Q

Explain why the value of the first ionisation energy increases from group 1 to group 2. [2]

A
  1. Higher nuclear charge;
  2. Electron being removed is in the same subshell
46
Q

Explain why the value of the first ionisation energy decreases from group 2 to group 3. [2]

A
  1. (The electron being removed) is from a new subshell;
  2. Which is further from the nucleus than the s-subshell
47
Q

State what is meant by the term relative atomic mass. [2]

A
  1. The (weighted) mean mass of an atom of an element;
  2. Compared to 1/12th of the mass of an atom of carbon-12
48
Q

Give the reason why, despite the difference in atomic structure, isotopes have the same chemical reactions. [1]

A
  1. They have the same electronic configuration
49
Q

State how the relative abundance of two isotopes can be found. [2]

A
  1. Compare the number of particles of each isotope detected;
  2. in a mass spectrometer
50
Q

Explain the term isotopes. [2]

A
  1. (atoms that have the) same number of protons;
  2. but different number of neutrons
51
Q

State what is meant by the terms ‘relative isotopic mass’ and ‘relative atomic mass’. [3]

A
  1. (relative isotopic mass refers to) the mass of an atom of that isotope;
  2. (relative atomic mass refers to) the weighted mean mass of an atom;
  3. relative to 1/12th of the mass of a carbon-12 atom
52
Q

Give a reason why the mass spectrometer must be operated under vacuum. [1]

A
  1. To prevent collisions with particles
53
Q

Give the meaning of the term ‘periodicity’.
Illustrate your answer by referring to the atomic radii of the Period 2 and Period 3 elements. [3]

A
  1. A trend of repeating (physical and chemical) properties (with increasing atomic number);
  2. Atomic radii decrease across the period;
  3. The atomic radius trend is repeated in period 3
54
Q

Explain the trend in melting temperatures across the elements of Period 2 in terms of their structure and bonding. [6]

A
  1. In Li and Be the bonding is metallic;
  2. Metallic bonding gets stronger as the number of delocalised electrons in a metal increases;
  3. C has a giant structure of atoms;
  4. A lot of energy is needed to break (strong) covalent bonds;
  5. N to Ne are simple molecules;
  6. Weak London forces (between molecules)
55
Q

Explain why two of the d-block elements within Period 4 (scandium and zinc) are not classified as transition metals.
You should include full electronic configurations where relevant. [6]

A
  1. (Transition metal) forms an ion with an incomplete d sub-shell;
  2. Scandium and zinc are not transition metals;
  3. Sc3+, 1s2 2s2 2p6 3s2 3p6
  4. Zn2+m 1s2 2s2 2p6 3s2 3p6 3d10;
  5. Sc3+ has an empty d sub-shell;
  6. Zn2+ has a full d sub-shell

Chemistry Exam Questions (3 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions