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CCEA AS Chemistry Units 1, 2 And 3 > (U1) Atomic Structure > Flashcards

(U1) Atomic Structure Flashcards

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1
Q

What is atomic number?

A

The number of protons in the nucleus of an atom

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2
Q

What is mass number?

A

The total number of protons and neutrons in the nucleus of an atom.

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3
Q

What is Relative Atomic Mass (RAM) (3)

A
  • The weighted mean mass
  • of an atom of an element
  • relative to 1/12 of the mass of an atom of carbon-12
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4
Q

What is Relative Isotopic Mass (RIM) (3)

A
  • The weighted mean mass
  • of an atom of an isotope of an element
  • relative to 1/12 of the mass of an atom of carbon-12
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5
Q

What are isotopes? (2)

A
  • Atoms which contain the same number of protons
  • but a different number of neutrons
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6
Q

What is Relative Molecular Mass (RMM) (3)

A
  • The weighted mean mass
  • of a molecule
  • relative to 1/12 the mass of an atom of carbon-12
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7
Q

What is Relative Formula Mass (RFM) (2)

A
  • The weighted mean mass of a species
  • relative to 1/12 the mass of an atom of carbon-12
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8
Q

What is a molecular ion?

A

2 or more atoms covalently bonded with an overall charge

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9
Q

What is First Ionisation Energy? (2)

A
  • The energy required to convert one mole of gaseous atoms
  • into gaseous ions with a single positive charge
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10
Q

What is Second Ionisation Energy? (2)

A
  • The energy required to convert one mole of gaseous ions with a single positive charge
  • into gaseous ions with a double positive charge
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11
Q

What is Third Ionisation Energy? (2)

A
  • The energy required to convert one mole of gaseous ions with a double positive charge
  • into gaseous ions with a triple positive charge
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12
Q

Using a mass spectrometer, how do you calculate relative atomic mass?

A

(percentage abundance) x (mass charge ratio) / 100

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13
Q

What is an orbital? (2)

A
  • a region within an atom that can hold up to 2 electrons
  • within opposite spins
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14
Q

What shape is an S-type orbital?

A

Spherical

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15
Q

What shape is a P-type orbital?

A

Dumbbell

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16
Q

State Hund’s Rule (2)

A
  • where electrons have choice between orbitals of equal energy
  • they will fill the orbitals singly as far as possible
17
Q

State the order of filling for energy levels (up to 9th)

A
  • 1s
  • 2s
  • 2p, 3s
  • 3p, 4s
  • 3d, 4p
18
Q

Why is 4s filled before 3d?

A

4s is closer to the nucleus than 3d

19
Q

In s, p, d, f notation, what do the letters represent?

How many electrons can s, p and d hold?

A
  • Sub shells
  • s holds 2, p holds 6, d holds 10
20
Q

What atom does the electronic configuration 1s2 represent?

What does this mean?

A
  • an atom of helium
  • the 1st ‘s’ sub shell is filled with 2 electrons
21
Q

State the general equation for 1st ionization energy

A

X(g) —> X+(g) + e-

22
Q

State the general equation for 2nd ionization energy

A

X+(g) —> X2+(g) + e-

23
Q

State the general equation for 3rd ionization energy

A

X2+(g) —> X3+(g) + e-

24
Q

Why does ionization energy usually increase across a period? (3)

A
  • increased nuclear charge
  • Decrease in atomic radius
  • same number of shielding electrons
25
Q

What is the electronic configuration of Cr (using noble gas notation)

Why is it irregular?

A
  • [Ar] 3d5 4s1
  • the 4s subshell loses an electron to the 3d subshell to form a half filled subshell (more stable)
26
Q

What is the electronic configuration of Cu (using noble gas notation)

Why is it irregular?

A
  • [Ar] 3d10 4s1
  • the 4s subshell loses an electron to the 3d subshell to form a filled subshell (stable)
27
Q

What 2 elements are exceptions to the rule for the filling of orbitals?

A
  • Cu
  • Cr
28
Q

Why does ionization energy not increase regularly across a period?

A
  • some have higher than expected energies due to filled and half filled subshells
  • transition metals increase marginally due to 3d subshell shielding 4s effectively
29
Q

What happens to 1st ionization energy down a group?

Why?

A
  • decreases
  • increase in atomic radius and more energy levels —> more shielding —> attraction weakens
30
Q

What are valence electrons?

A

Outer electrons

31
Q

How many orbitals are in a:

  • p-subshell
  • s-subshell
  • d-subshell
A
  • p = 3 orbitals
  • s = 1 orbital
  • d = 5 orbitals
32
Q

State and explain the general trend in the first ionization energy across the second period.

Identify and explain any exceptions

(7)

A
  • 1st ionization energy increases left to right
  • due to:
    1. increase in nuclear charge
    2. Decrease in distance between nucleus and outer electrons
    3. Same number of shielding electrons
  • exceptions: Be and N
  • Be —> full outer subshell —> stability —> more energy needed
  • N —> half full outer subshell —> stability —> more energy needed
33
Q

When forming cations, which orbital are electrons removed from first in Cr?

Why?

A
  • The 4s subshell
  • 4s is further away from the nucleus than 3d
34
Q

What are the relative masses and charges of:

  1. Protons
  2. Neutrons
  3. Electrons
A
  1. Protons
  • mass = 1
  • charge = +1
  1. Neutrons
  • mass = 1
  • charge = 0
  1. Electrons
  • mass = 1/1840
  • charge = -1
35
Q

How do you calculate relative atomic mass from mass spectra?

(equation)

A

RAM = percentage abundance (%) x mass charge ratio (Z) / 100

36
Q

On the periodic table, where is the:

  • s-block
  • p-block
  • d-block
  • f-block

Why are they called this?

A
  1. S-block = groups 1 and 2
  2. P-block = groups 5 to 8
  3. D-block = transition metals
  4. F-block = lanthanum series

Name comes from their outermost subshell

37
Q

In box notation, how do you represent electrons?

A

⥮ one arrow up, one arrow down - representing 2 electrons in opposite spins (within an orbital) ⥮

38
Q

What does a large difference (compared to 3rd ionisation energy etc.) in 2nd ionisation energy confirm?

Why? (2)

A

The presence of energy levels:

  • first electron lost easily as it is in a shell furthest from the nucleus (greater atomic radius) and separated by greater shielding, therefore electrostatic attraction is weakest
  • 2nd electron closer to nucleus and requires more energy to overcome stronger electrostatic attraction
39
Q

What is the affect of ionisation on effective nuclear charge?

A

Increased ionisation = greater effective nuclear charge

(greater ratio of protons to electrons)

CCEA AS Chemistry Units 1, 2 And 3 (22 decks)

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