Chapter 1 - Atomic Structure Flashcards Preview

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Flashcards in Chapter 1 - Atomic Structure Deck (25):
1

How do particles behave when placed in an electric field? (3)

1) protons&cations: deflected to negative plate
2) electrons&anions: deflected to positive plate
3) neutrons: pass straight through

2

How is the angle of deflection calculated?

it is proportionate to q(charge)/m. Directly proportional to size of charge and inversely proportional to mass

3

What is the principal quantum number(n)?

It is the main energy level of an electron and also indicates the relative size of the orbital and therefore the relative distance of the electron from the nucleus.
The larger the value of n, the higher the energy level and the further the electron is from the nucleus.

4

What is the maximum number of electrons that can occupy each principal quantum shell?

2n^2

5

What is a subshell?

It is subdivided from a principal quantum shell. The number of subshells in a principal quantum shell is the same as its principal quantum number.

6

What are the 4 types of sub shells and how do they differ in energy level?

s, p, d, f, with s closest to nucleus. s

7

What is an atomic orbital?

It is a region of space with a 90% probability of finding an electron, and has a specific energy associated with them.

8

What are 3 properties of orbitals?

1) It can contain up to a maximum of 2 electrons which are of opposite spins.
2) Orbitals in the same sub shell has the same energy level (degenerate) but different orientations in space.
3) size of orbitals increase as quantum number increases

9

How are electrons distributed in various orbitals? (4 rules)

Rules:
1) Orbitals with the lowest energy is always filled first. Electrons only enter higher energy levels after lower energy levels have been filled.
2) in the same sub shell, with orbitals of the same energy, each orbital must be singly occupied before electrons are paired.
3) 2 electrons in the same orbital must have opposite spins (except chromium and copper)
4) when removing electrons (such as forming cations), remove electrons from the outermost (highest energy) sub shell first.

10

Define transition metals.

A transition metal is a d-block (middle block) element that forms some compounds containing its ion with an incomplete d-shell.

11

Define effective nuclear charge(Zeff).

It is the net nuclear charge experienced by an outer electron (attraction of valence electron towards nucleus).

12

What is the effective nuclear charge dependent on? (2)

1) Nuclear charge(Z): number of protons. The larger the number of protons, the greater the nuclear charge and the greater the attraction for the outermost electrons by the nucleus.
2) Shielding effect(S): Presence of inner-shell electrons reduced the electrostatic forces of attraction between the outermost electrons and nucleus.
Zeff=Z-S

13

What are 2 properties of shielding effect?

1) electrons in the same quantum shell do exert some shielding effect on each other, but it is small compared to inner core electrons.
2) shielding extent increases: s>p>d>f

14

What is the general trend of atomic radius down the group?

Atomic radius decreases down the group.
- number of quantum shell increases (as additional electrons are added to a new quantum shell). The outermost electrons are further away from the nucleus, hence increasing atomic radii
- Although nuclear charge also increases down the group due to increase in number of protons in the nucleus, the number of quantum shells is a more important factor to consider.

15

What is the general trend of atomic radius across period 2 and 3?

Atomic radius decreases down the group.
- nuclear charge increases due to increase in number of protons in nucleus
- shielding effect remains relatively constant (as electrons are added to the same outermost shell)
- outermost shell are pulled closer to the nucleus and hence decreasing atomic radii

16

What is the general trend of atomic radius across first row of transition elements?

Atomic radius is relatively invariant (constant).
- nuclear charge increases due to increasing number of protons
- electrons are added to the inner 3d sub shell as we go across the period, which contributes to the increase in shielding effect
- the shielding effect increases and nullifies (deprive effectiveness), to a considerable extent, the influence of each additional proton in the nucleus.
- effective nuclear charge thus remains relatively constant

17

What is the general trend of ionic radius?

Across an isoelectronic series (ions with same number of electrons), ionic size decreases.
- nuclear charge increases due to increasing number of protons
- shielding effect remains the same due to same number of electrons
- effective nuclear charge increases. stronger attraction between outermost electrons and nucleus decreases ionic size
- ps: there is a sharp increase in ionic radius from cationic series of Na+ to Al3+ to the anionic series of P3- to Cl- as the anions have one more quantum shell of electrons than cations.

18

Define 1st ionisation energy. (and nth ionisation energy)

It is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of unipositively charged gaseous ions.
It is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms with (n-1)+ charge to form 1 mole of gaseous ions with n+ charge.

19

What are 2 factors affecting ionisation energy of an atom?

1) Number of quantum shells: the larger the number of quantum shells, the further the electrons is from the nucleus and hence experiences weaker nuclear attraction and is easier to remove, thus the lower the ionisation energy.
2) Effective nuclear charge: the higher the effective nuclear charge, the stronger the attractive forces between the nucleus and the harder for electrons to the removed, thud the greater the ionisation energy.

20

What is the trend of successive ionisation energy in the same element?

For the same element, ionisation energy increases as n increases.
When an electron is removed from a neutral atom, the number of protons that exert an attraction for the remaining electrons remain the same. However, the shielding effect among the remaining electrons in the outermost shell is reduced since there is now one less electron. Hence, effective nuclear charge increases. More energy is needed to remove another electron from the more positively charged ion, hence a higher ionisation energy.

21

How can we determine the shell number or electronic configuration of an element through its successive ionisation energies?

gradual increase: same shell
small increase: next sub shell
large increase: next quantum shell
越来越里面

22

What is the trend of first ionisation energies of different elements down the group?

It decreases down the group.
Although there is an increase in nuclear charge down the group, the increase in number of quantum shells of electrons means that the outermost electrons are further from the nucleus. Therefore, the electrostatic forces of attraction between the nucleus and the outermost electron are weaker and less energy is required to remove this electron.

23

What is the trend of first ionisation energies of different elements across periods 2 and 3?

It increases across a period.
Across the period, nuclear charge increases and shielding effect remains relatively constant, hence increasing effective nuclear charge (same reason as atomic radius). The electrostatic forces of attraction between the outermost electrons and the nucleus becomes stronger, so more energy is required to remove the outermost electron.

24

What are 2 anomalies (反常) of the trend of first ionisation energies of different elements across periods 2 and 3?

1) there is a small dip between group 2 and 13 electrons as Al has a lower 1st ionisation energy than Mg. The 3p sub shell of Al is further away from the nucleus than the 3 sub shell of Mg. There is a weaker attraction between the nucleus and the outermost electron. Hence, less energy is needed to remove the 3p electron from Al, resulting in a lower ionisation energy for Al.
2) There is a small dip between group 15 and 16 elements as S has a lower 1st ionisation energy than P. This is because all the 3p electrons in P are unpaired, while 2 of the 3p electrons are paired in S. There is some inter-electronic repulsion between the paired electrons in the 3p sub shell of s, thus less energy than expected is required to remove one of these paired electrons from S.

25

What is the trend of first ionisation energies of different elements across transition elements?

First ionisation energy remains relatively invariant.
1st ionisation energy involves the removal of a 4s electron.
- The nuclear charge increases due to the increasing number of protons
- additional electrons are added to the inner 3d sub shell, which contributes to the shielding effect.
- the shielding effect increases, thereby nullifying, to a considerable extent, the influence of each additional proton in the nucleus.
- effective nuclear charge thus remains almost constant. Energy required to remove the outermost electron in each succeeding element remains relatively invariant.