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IB Chemistry > 1. > Flashcards

1. Flashcards

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1
Q

heterogenous

A

substance where each component exists in clumps. You can also see induvidual components. e.g. oil and water

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2
Q

homogenous

A

Substance where each component is evenly distributed. Difficult to identify the induvidual components. e.g. coffee

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3
Q

distillation

A

seperates solvent from a solution

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4
Q

filtration

A

insoluble solid from liquid

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5
Q

crystallisation

A

Seperating solute from a solution

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6
Q

chromatography

A

e.g. ink or coloured compounds

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7
Q

fractional distillation

A

seperates a mixture using boiling points

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8
Q

Solvation

A

Seperating insoluble + soluble solids using a solvent

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9
Q

Limonene

A

An organic compound found in the rinds of citrus fruits.
- strong aroma
- has potential health benefits but no groundbreaking evidence

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10
Q

Qualitive

A

Quality, anything that can be detected by senses

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11
Q

Quantitive

A

quantity

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12
Q

Gas to solid

A

Deposition

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13
Q

Solid to gas

A

Sublimation

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14
Q

Are kelvin and Celsius the same kind of scale?

A

Yes kelvin starts a 0 and goes up to 373.15 where water boils (373.15 diff)

Celsius starts at -273.15 and goes to 100 where water boils (373.15 diff)

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15
Q

What are the limitations of the particle model?

A
  • Particles aren’t moving
  • only pure substances taken into account
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16
Q

How do nuclei of atoms differ?

A
  • Different number of protons.
  • elements have same proton number but different nucleon number which is an isotope.
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17
Q

What did Democritus propose?

A

Proposed all matter was made up of smaller unites named “atomos” meaning invisible

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18
Q

What did JJ Thompson discover?

A

In 1906 he discovered the electron thinking it was in the nucleus with protons.

Made a model known as the plum pudding model

19
Q

What was the gold foil experiment and what were the findings?

A

The gold foil experiment was where alpha particles (He2+ particles) were fired into gold foil and there was a large screen to record findings. Very rarely almost 1 in 20,000 the particle would repel or even come back to where the alpha particle were fired. Happened in 1911

WHAT DID THESE FINDINGS SUGGEST?

  1. Most of the atom is empty space.
  2. Positively charged centre with electrons revolving around the nucleus
  3. Size of nucleus is small compared to size of atom
20
Q

What did Niels Bohr discover?

A

Electrons move around a nucleus but only in prescribed orbits (e.g. 2,8 …) like the solar system the electrons orbit the nucleus like we orbit the sun.

21
Q

Nuclear symbol

A

E.g. Magnesium

A *
MG
Z **

  • atomic mass (protons + neutrons)

** atomic number (number of protons)

22
Q

Isotope

A

Same number of protons but different number of neutrons.

23
Q

Electromagnetic spectrum

A

Radio waves
Microwaves
Infrared
Visible light
UV light
X-ray
Gamma Rays

Radio waves to gamma rays frequency increases

Gamma rays to radio waves wavelength increases

24
Q

What is the visible spectrum?

A

Section of electromagnetic spectrum which is visible to the human eye

25
Q

Emission spectrum

A

If a pure gaseous element is subjected to a high voltage under reduced pressure, the gas will emit a certain characteristic colour of light. E.g. sodium emits yellow

26
Q

Quantum’s of radiation

A

A quantum of radiation is called a photon. It has a specific wavelength that relates to a discrete value of energy. Each line on an emission spectrum is releasing a specific photon of radiation. This is called quantilisation.

27
Q

What is a quantum number?

A

E.g n=1

28
Q

Unlike IGCSE what are electron shells called?

A

Energy levels

29
Q

Emission spectra for hydrogen

A

Provides evidence for existence of electrons in discrete energy levels, which converge at higher energies

nah to n=1 is the Lyman series and emits ultraviolet

nh to n=2 is the Balmer series which emits visible light

nh to n=3 is the Paschen series which emits infrared

30
Q

What is an energy level?

A

Each assigned a quantum number, n. Each energy level can hold 2n^2 electrons.

31
Q

What are the sub-levels called?

A

S orbital } most likely in exam
P orbital } most likely in exam
D orbital
F orbital

32
Q

What is an atomic orbital?

A

Each orbital contains fixed orbitals which are regions of space where there is a high probability chance for finding an electron. Each orbital has defined energy state.

33
Q

How many electrons can each orbital hold?

A

2 electrons with opposite spins (one spin up electron and one spin down electron)

34
Q

What are the electron configuration rules?

A
  1. Aufbau principle - electrons must fill up the lowest energy orbital first. Usually the case.
  2. Pauli exclusion principle - any orbital can hold a maximum of two electrons which have opposite spins.
  3. Hunds rule - any electrons fill all orbitals in a sub level before occupying them in pairs.
35
Q

What are the features of a transition metal?

A
  • has valence electrons in two shells
  • form coloured compounds
  • metals with different oxidation states
  • can be catalysts
  • can form complex ions
36
Q

Why do all transition metals form +2?

A

All have 2 electrons in a block, 1s^2. These electrons in a block are taken first as they have the highest quantum number.

37
Q

Why does copper only have one electron in the 1s block?

A

For stability reasons meaning there is 10 instead of 9 in the d block.

38
Q

First ionisation energy

A

The energy required to remove one mol of electrons from one mol of gaseous atoms to form one mol of gaseous +1 ions under standard conditions.

39
Q

How to find the first ionisation energy?

A

hc/wavelength

ANS x 6.02x10^23

ANS/1000

40
Q

How to find first ionisation energy?

A

If you know the wavelength and if it’s in nm (if not convert it)

hc/wavelength

ANS x 6.02x10^23

ANS/1000

41
Q

How to find first ionisation energy?

A

If you know the wavelength and if it’s in nm (if not convert it)

hc/wavelength

ANS x 6.02x10^23

ANS/1000

42
Q

What factors impact ionisation energy?

A
  1. Nuclear charge - is the attraction of the nucleus to the electrons. The greater the number of protons the greater the nuclear charge. Increases ionisation energy as you go down the periodic table.
  2. Electron shielding - is electron-electron repulsion between energy levels. Full energy levels ‘shield’ valence electrons from the nuclear attraction of the nucleus. Decreases ionisation energy as you go down the periodic table
  3. Electron spin repulsion - is when an electron is added to an atomic orbital in a half-filled sub-level. Decreases initiation energy as you go down the group.
  4. Atomic radius - is the distance between the nucleus and the valence electrons. The larger the gap the less ionisation required meaning it decreases as you go down the periodic table.
43
Q

Why is nitrogen require more first ionisation energy than oxygen.

A

Because nitrogen’s half filled configuration makes it more stable. The higher the stability the more first ionisation energy required to remove one mol of electrons.

44
Q

How does the first ionisation energy change along the groups/periods?

A

First ionisation energy decreases down a group. Across a period it increases with dips in groups 13 and 16.

IB Chemistry (3 decks)

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