SA 2 Flashcards

1
Q

In the line emission spectrum of hydrogen, which of the following electron transitions would not appear?
a. n = 3 to n = 1
b. n = 3 to n = 4
c. n = 4 to n = 1
d. n = 3 to n = 2

A

b. n = 3 to n = 4

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2
Q

In which energy level can we see visible lines

A

n = 2

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3
Q

What is meant by a ‘metallic character’?

A

Defined as how easily an atom can lose an electron

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4
Q

What did Bohr conclude about the structure of the atom from analysis of this spectrum?

A

Electrons move around the nucleus of the atom in certain fixed paths.

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5
Q

Which of the following has the greatest metallic character?

a. Cl
b. Na
c. Al
d. C

A

b. Na

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6
Q

What is the approximate bond angle between Cl–N–Cl in NCl3?

A

109.5°

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7
Q

What intermolecular molecules does SO2 have?

A

dipole-dipole and dispersion forces

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8
Q

1s2 2s2 2p6 3s2 3p6 3d4 5s1

(i) Identify the atom.

(ii) Write the ground state electron configuration of the atom.

(iii) Explain what is meant by the term ‘ground state’ for the electron configuration.

A

i. Vanadium
ii. 1s2 2s2 2p6 3s2 3p6 4s2 3d3
iii. ground state = the lowest possible electron configuration (lowest energy level)

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9
Q

Draw the Lewis structure and Identify the molecular shape of:

(i) SiF4
(ii) SF4

A

google lol

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10
Q

What is the main concept of the Bohr model?

A

Electrons occupy definite orbits and each orbit has a specific amount of energy - aka quantised. Orbits closer to the nucleus have smaller amounts of energy, to go to a higher energy level, it needs to absorb energy, and vice versa (to go to a lower energy level, emit energy). The amount of energy absorbed or emitted is fixed.

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11
Q

How do atoms gain energy?

A

One way is to pass an electric current through an enclosed sample of a gas. Another way is to supply heat energy using a flame (flame tests). Then, the light emitted can be passed through a prism to separate the different wavelengths of light. Only discrete lines of certain wavelengths appear, the spectrum is called a line emission spectrum.

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12
Q

What was the problem with the Bohr model?

A

That it fails for any multi-electron system.
The Bohr model is only successful for atoms that have a single electron, energy levels containing more than one atom cannot be calculated successfully.

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13
Q

What is the quantum mechanical model?

A

A model made by Schrödinger, which accurately described the behaviour of the electron in a hydrogen atom. (note that its not a ‘model’, but a s p f d thing)

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14
Q

What is the Pauli Exclusion Principle?

A

That no two electrons can have the same set of quantum numbers - every orbital has a max of 3 electrons, and can be unoccupied, occupied by 1, or occupied by 2.

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15
Q

What is the difference between the Bohr model and the Quantum mechanical model?

A

In the Bohr model, the pathway of the electron was fixed into circular orbits. The quantum mechanical model states that the electrons do not orbit the nucleus but instead travels in specific patterns referred to as electron clouds.

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16
Q

What is the Aufbau Principle?

A

Boxes. That electrons occupy the lowest energy to the highest. s then p then d then f.

17
Q

What is Hund’s Rule?

A

That orbitals of equal energy are occupied by one electron before any orbital is occupied by a second electron, maximizing the number of unpaired electrons. Boxes.

18
Q

What elements are the exceptions to the ground state configurations?

A

Cr = 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Cu = 1s2 2s2 2p6 3s2 3p6 4s1 3d10

19
Q

Write 1s2 2s2 2p6 3p1 in its excited state

A

1s2 2s2 2p6 4p1

20
Q

What are these called?
Group 1, 2, 17, 18

A

Group 1 are known as the alkali metals
Group 2 is also known as alkali earth metals
Group 17 is also known as the halogens
Group 18 is also known as the noble gases

21
Q

What is an isoelectronic species?

Compare the sized of:
O2-, F-, Ne, Na+ and Mg2+

A

Elements that have the same number of electrons (Li+ and Be2+), they will have the same amount of shielding but different number of nuclear charge (more nuclear charge = smaller)

Mg2+ < Na+ < Ne < F- < O2-

22
Q

What is ionisation energy?
State the first ionisation energy equation

A

The energy required to remove an electron from an atom in its ground state
X(g) → X+(g)+ e-

23
Q

What is electronegativity and what is its difference with electron affinity?

A

The ability of an atom to attract electrons in a bond when the atom is part of a compound.

  • different from electron affinity because electron affinity is the actual amount of energy released when an atom gains an electron. Electronegativity is not measured in energy units, but is a relative scale, where the elements are compared to each another.
24
Q

What is the octet rule?

A

The losing or gaining of an electron to achieve the lowest possible noble gas configuration.

25
Q

What are the exceptions to the octet rule?

A
  1. Elements starting period 3 (d orbital is available)
  2. Coordinate covalent compounds (1 atom providing both electrons in a shared pair, eg: ammonia and H2)
  3. In some compounds, the number of electrons in the core of a stable molecule is less than 8 (Be - doesn’t form ionic bonds due to small size and high ionisation energy)
26
Q

What is the Valence-shell Electron Repulsion Theory (VSEPR)?

A

States that the molecule will adjust its shape so that electrons stay as far as possible

27
Q

What are dipole bonds and what are its 2 types?

A

Secondary bonds where the electrons are not evenly shared, only occur between polar molecules
1. ion-dipole bonds (strength depends on size of ion, charge of ion, size of dipole)
2. dipole-induced dipole (weak attraction resulting from a polar molecule from a polar molecule influencing a dipole in an atom or in a non-polar molecule - O2 getting repelled by H2O due to H2O having more electronegativity)

28
Q

What can H bond with to make a hydrogen bind?

A

O, N, F (H bonds are polar)

29
Q

Which one of the following best accounts for the general increase in ionisation energy across a row of the Periodic Table?

a. The increasing number of electrons in the atoms of the elements across a row.

b. The decreasing distance between the nucleus and the outer electrons in the atoms of the elements across a row.

c. The increasing effective nuclear charge within the atoms of the elements across a row.

d. The increasing number of neutrons in the atoms of the elements across a row.

A

c. The increasing effective nuclear charge within the atoms of the elements across a row.

30
Q

Which of the following bonds is the most polar?

a. H – H

b. F – F

c. H – I

d. H – F

A
  1. Polarity relates to electronegativity
    d. H – F
31
Q

Which of the following compounds contains only covalent primary bonds?

a. CH4

b. K

c. NaCl

d. Na2CO3

A

a. CH4

(K is metallic, NaCl is ionic, Na2CO3 has both ionic and covalent compounds)

32
Q

Which statement best explains why oxygen cannot bond with fluorine to form OF6, while sulfur can bond with fluorine to form SF6?

a. Sulfur has more valence electrons than oxygen.

b. The 3d orbitals in sulfur can be used in bonding.

c. There are less valence electrons in oxygen.

d. Oxygen is a larger atom.

A

b. The 3d orbitals in sulfur can be used in bonding.

33
Q

How many lone pairs are around the boron in BH3 (B is the central atom)?

A

0

34
Q

What is the molecular shape of NCl3?

A
  1. look at lone pairs
    trigonal pyramidal
35
Q

What is the approximate bond angle between F-O-F in OF2?

A
  1. look at lone pairs
    109.5°
36
Q

Which of the following bonds is the most polar?

a. H – H

b. F – F

c. H – I

d. H – F

A
  1. look at elecneg
    d. H – F
37
Q

Which of the following fluorides is polar?

a. NF3

b. PF5

c. CF4

d. SiF4

A
  1. look at lone pairs