Module 2.1 Atoms and Reactions Flashcards Preview

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Flashcards in Module 2.1 Atoms and Reactions Deck (87)
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1

15th C - The Greek atom

Greek philosopher Democritus developed the first idea of the atom

Suggested you could divide matter only a certain number of times - "átomos" meaning "indivisible"

2

Early 1800s - Dalton's atomic theory

Atoms are tiny particles that make up elements

Atoms cannot be divided

All atoms of a given element are the same

Atoms of one element are different from those of every other element

Dalton used his own symbols to represent atoms

Dalton developed the first table of atomic masses

3

1897-1906 - Joseph John Thomson

He discovers electrons

Discovered "cathode rays" had a negative charge, could be deflected by a magnet and an electric field and had a v v small mass

Cathode rays were electrons

The idea that an atom could not be split any further (proposed by the ancient Greeks and Dalton) had been disproved

4

1897-1906 - Joseph John Thomson

Proposed that atoms were made up of negative electrons moving around a "sea" of positive charge - "plum pudding model"

He thought that the overall negative charge = the overall positive charge - an atom is neutral with no overall charge

5

1909-1911 Ernest Rutherford

Gold leaf experiment:
Directed a-particles towards a sheet of very thin gold foil. They measured any deflection of the particles.

What the results showed:
Most particles weren't deflected at all

A small percentage of particles deflected at large angles

Few particles were deflected back towards the source

His new model (1911):
The positive charge of an atom and most of the mass are concentrated in the nucleus, in the centre

Negative electrons orbit the nucleus - just like how planets orbit the sun

The + and - charges must balance

Rutherford had disproved the plum pudding model - the expected result was that the a-particles would pass through undisturbed but the observed results were that a few particles were deflected - indicating a small concentrated positive charge (the nucleus)

6

1913 - Niels Bohr and Henry Moseley

Bohr thought electrons could only follow certain paths otherwise they would spiral into the nucleus - "planetary atom" - electrons orbited the central nuclear "sun" in "shells"

Bohr's model helped explain some periodic properties e.g. spectra lines seen in the emission spectra, the energy of electrons at different distances from the nucleus

Henry Moseley discovered the link between x-ray frequencies and an element's atomic number - however, he couldn't explain this at the time

7

1918- Rutherford

Discovers the proton

Could explain Moseley's finding that an atom's atomic number was linked to x-ray frequencies

We now know that atomic no. = no. of protons in the atom

8

1923-1926

Louis de Broglie suggested that particles could have the nature of both a wave and a particle

Erwin Schrödinger suggested that an electron had wave-like properties in an atom. He also introduced the idea of atomic orbitals

9

1932- James Chadwick

Discovers neutrons

Observed a new type of radiation emitted from some elements

Uncharged particles - approx. same mass as a proton

These were neutrons

10

Modern day

Now thought that protons and neutrons are made up of even smaller particles called quarks

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Current model of the atom

Protons and neutrons in the nucleus

Nucleus = centre of atom

Electrons orbit the nucleus in "shells"

Nucleus = absolutely tiny

Nucleus = extremely dense, accounts for almost all of the atom's mass

Most of an atom is empty space

12

Relative mass and charge of the sub-atomic particles

Proton:
Mass = 1
Charge = +1

Neutron:
Mass = 1
Charge = 0

Electron:
Mass = 1/2000
Charge = -1

13

Isotopes

Different masses

Same no. of protons

Same no. of electrons

Different number of neutrons

14

Atomic number

Proton number

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Mass number

Protons + neutrons

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Relative molecular mass, Mr

Add together the relative atomic masses of each atom making up a molecule

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Relative formula mass

Add together the relative atomic masses of each atom making up the formula unit

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What is mass spectrometry used for?

Identify an unknown compound

Find the relative abundance of each isotope of an element

Determine structural information about molecules

19

A mass spectrum shows

Relative or percentage abundance on the y-axis and mass-to-charge ratio on the x-axis

20

How is mass-to-charge ratio in all mass spectra shown as

m/z

m= mass
z= the charge of the ion (usually 1)

21

Atoms of metals in groups 1-13

Lose electrons

Form positive ions with the electron configuration of the previous noble gas in the periodic table

22

Atoms of non-metals in groups 15-17

Gain electrons

Form negative ions with the electron configuration of the next noble gas in the periodic table

23

Why do atoms of Be, B, C and Si not usually form ions?

It requires too much energy to transfer the outer shell electrons to form ions

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Ammonium

NH4+

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Hydroxide

OH-

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Nitrate

NO3-

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Carbonate

CO3^2-

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Sulphate

SO4^2-

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Avogadro's constant

6.02 times 10^23 mol^-1

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Number of moles

Mass divided by molar mass