atomic estruture Flashcards
(76 cards)
1
Q
A
2
Q
What are the smallest parts of an element that can take place in chemical reactions?
A
Atoms
Atoms are composed of subatomic particles: protons, neutrons, and electrons.
3
Q
What is the structure of an atom primarily composed of?
A
Empty space around a dense nucleus
The nucleus contains protons and neutrons.
4
Q
What charge does the nucleus of an atom have?
A
Overall positive charge
5
Q
What are the charges of protons and neutrons?
A
Protons: +1, Neutrons: 0
6
Q
Where are negatively charged electrons found in an atom?
A
In orbitals around the nucleus
7
Q
What are subatomic particles?
A
Protons, neutrons, and electrons that make up an atom
8
Q
How are the masses and charges of subatomic particles expressed?
A
Relative atomic masses and relative atomic charges
9
Q
What is the relative mass of protons and neutrons?
A
1
10
Q
What is the relative mass of electrons compared to protons and neutrons?
A
1 over 1836
11
Q
What is the charge of a single electron in coulombs?
A
-1.602 x 10^-19
12
Q
What is the charge of a proton in coulombs?
A
+1.602 x 10^-19
13
Q
What holds an atom together?
A
Electrostatic attraction between the nucleus and electrons
14
Q
How do protons, neutrons, and electrons behave in an electric field?
A
Electrons are deflected towards positive plates; protons are deflected towards negative plates; neutrons are not deflected.
15
Q
What is the atomic number (Z)?
A
Number of protons in the nucleus of an atom
16
Q
What does the mass number (A) represent?
A
Total number of protons + neutrons in the nucleus
17
Q
How can you calculate the number of neutrons in an atom?
A
Mass number - atomic number
18
Q
What happens to an atom when it gains or loses electrons?
A
It becomes an ion
19
Q
What is the mass number of an atom with 6 protons and 6 neutrons?
A
12
20
Q
In a neutral atom, how many electrons does it have compared to protons?
A
The same number
21
Q
What is the atomic radius?
A
Half the distance between the two nuclei of two covalently bonded atoms
22
Q
How does atomic radius change across a period in the Periodic Table?
A
Decreases
23
Q
How does atomic radius change down a group in the Periodic Table?
A
Increases
24
Q
What is an isotope?
A
Atoms of the same element with the same number of protons but different number of neutrons
25
What are the chemical properties of isotopes?
Similar chemical properties
26
What are the physical properties of isotopes?
Different physical properties such as mass and density
27
What is the arrangement of electrons in an atom called?
Electronic configuration
28
What are the principal quantum numbers used for?
Numbering the energy levels or quantum shells
29
What is the maximum number of electrons that can occupy the first principal quantum shell (n=1)?
2 electrons
30
What is the order of energy levels in sub-shells?
s < p < d
31
What is the maximum number of electrons that can occupy an atomic orbital?
2 electrons
32
What are the energy levels of the quantum shells?
Energy levels correspond to the principal quantum number n. At n = 3, energy levels of sub-shells start to overlap.
## Footnote
The principal quantum number indicates the energy level of a particular shell.
33
What defines an orbital in atomic structure?
An orbital is a region where electrons can be found at specific energy levels, each orbital can hold a maximum of two electrons.
## Footnote
Orbitals exist at specific energy levels and electrons can only be found at these specific levels.
34
How many orbitals are in each type of sub-shell?
* s: 1 orbital (2 electrons)
* p: 3 orbitals (6 electrons)
* d: 5 orbitals (10 electrons)
* f: 7 orbitals (14 electrons)
## Footnote
The maximum number of electrons is calculated as 2 times the number of orbitals.
35
What is the shape of s and p orbitals?
* s orbitals: spherical
* p orbitals: dumbbell-shaped
## Footnote
The shapes of the orbitals differ, with s orbitals being spherical and p orbitals oriented along the x, y, and z axes.
36
What is the ground state of an atom?
The ground state is the most stable electronic configuration with the lowest energy, achieved by filling sub-shells in order of increasing energy.
## Footnote
The lowest energy sub-shells are filled first, starting with 1s.
37
What is the arrangement of electrons in shells and sub-shells?
* Principal quantum number (n): 1, 2, 3, 4
* Sub-shells: s, p, d, f
* Orbitals per sub-shell: s: 1, p: 3, d: 5, f: 7
* Electrons per sub-shell: s: 2, p: 6, d: 10, f: 14
* Electrons per shell: 1: 2, 2: 8, 3: 18, 4: 32
## Footnote
This summarizes the organization of electrons in atomic structure.
38
How does the energy of sub-shells increase?
The energy of sub-shells increases in the order: s < p < d < f, with the exception of 3d being higher in energy than 4s.
## Footnote
This is why the 4s orbital is filled before the 3d orbital.
39
What is the shape and size behavior of s orbitals?
s orbitals are spherical and their size increases with increasing shell number.
## Footnote
The s orbital of the third quantum shell (n = 3) is larger than that of the first (n = 1).
40
What is the configuration of electrons in p orbitals?
The p orbitals are dumbbell-shaped and every shell has three p orbitals except for the first one (n = 1).
## Footnote
The lobes of p orbitals increase in size with higher shell numbers.
41
What is an electron configuration?
An electron configuration describes the distribution of electrons among the shells, sub-shells, and orbitals of an atom.
## Footnote
The sub-shells are filled in order of increasing energy.
42
What is the significance of electron spin?
Electrons can spin in either clockwise or anticlockwise directions, represented by their direction in diagrams.
## Footnote
Electrons with similar spins repel each other, leading to their arrangement in separate orbitals.
43
What defines a free radical?
A free radical is a species with one or more unpaired electrons, represented by a dot in its configuration.
## Footnote
For example, a chlorine free radical has an unpaired electron in one of its p orbitals.
44
How are ions formed in atomic structure?
Ions are formed when atoms lose or gain electrons, with negative ions formed by adding electrons and positive ions by removing them.
## Footnote
Transition metals fill the 4s before the 3d but lose electrons from the 4s first.
45
What are the four main blocks of the Periodic Table?
* s block: valence electrons in s orbital
* p block: valence electrons in p orbital
* d block: valence electrons in d orbital
* f block: valence electrons in f orbital
## Footnote
Elements are categorized based on their outer shell electron configuration.
46
What are the exceptions in electron configurations for chromium and copper?
Chromium: [Ar] 4s1 3d5; Copper: [Ar] 4s1 3d10.
## Footnote
These configurations are stable due to their energy arrangements.
47
Define ionisation energy.
Ionisation energy (IE) is the energy required to remove one mole of electrons from one mole of gaseous atoms to form ions.
## Footnote
Measured under standard conditions (298 K, 101 kPa) and always positive.
48
What factors affect the size of ionisation energy?
* Size of nuclear charge
* Distance of outer electrons from the nucleus
* Shielding effect of inner electrons
* Spin-pair repulsion
## Footnote
These factors influence the energy required to remove electrons.
49
What happens to ionisation energy as the outer electron shell is farther from the nucleus?
Ionisation energy decreases due to weaker nuclear attraction.
50
What is the shielding effect?
Electrons in full inner shells repel outer shell electrons, reducing the effect of nuclear charge.
51
How does spin-pair repulsion affect ionisation energy?
Electrons in the same orbital repel each other, making it easier to remove an electron, lowering ionisation energy.
52
What trend is observed in ionisation energy across a period?
Ionisation energy increases across a period.
53
What factors contribute to the increase in ionisation energy across a period?
* Increase in nuclear charge
* Decrease in atomic radius
* Constant shielding by inner electrons
54
What causes the rapid decrease in ionisation energy between the last element in one period and the first element in the next?
* Increased distance from nucleus to outer electrons
* Increased shielding by inner electrons
55
What is the first ionisation energy of Beryllium?
900 kJ mol-1
56
What is the first ionisation energy of Boron?
800 kJ mol-1
57
What is the first ionisation energy of Nitrogen?
1400 kJ mol-1
58
What is the first ionisation energy of Oxygen?
1310 kJ mol-1
59
What causes the slight decrease in ionisation energy between nitrogen and oxygen?
Spin-pair repulsion in the 2px orbital of oxygen.
60
What trend is observed in ionisation energy down a group?
Ionisation energy decreases down a group.
61
What factors contribute to the decrease in ionisation energy down a group?
* Increased nuclear charge
* Increased atomic radius
* Increased shielding by inner electrons
62
What happens to the distance between the nucleus and outer electrons as you descend a group?
The distance increases.
63
What happens to the shielding effect as you move down a group?
The shielding effect increases.
64
What is the trend of ionisation energy when comparing full subshells to principal quantum shells closer to the nucleus?
It is easier to remove electrons from full subshells than from principal quantum shells closer to the nucleus.
65
What is the general trend of successive ionisation energies of an element?
Successive ionisation energies increase.
66
Why does the first ionisation energy have a low value?
The first electron is easily removed due to spin-pair repulsion.
67
What is the first ionisation energy of Calcium?
590 kJ mol-1
68
What is the second ionisation energy of Calcium?
1150 kJ mol-1
69
What is the third ionisation energy of Calcium?
4940 kJ mol-1
70
What is the fourth ionisation energy of Calcium?
6480 kJ mol-1
71
What factors affect the magnitude of ionisation energy?
* Nuclear charge
* Shielding
* Atomic/ionic radius
* Spin-pair repulsion
72
True or False: Less energy is required to remove an outer shell electron from an atom with high nuclear charge.
False
73
What happens to the attractive forces between the outer electron(s) and the nucleus with increased shielding effects?
The attractive forces weaken.
74
Fill in the blank: The larger the atomic radius, the _______ the ionisation energy.
lower
75
What is the relationship between nuclear charge and ionisation energy?
Higher nuclear charge results in higher ionisation energy.
76
What effect does spin-pair repulsion have on ionisation energy?
It decreases ionisation energy.
