Module 2.1 Atoms and Reactions Flashcards
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”
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
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
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
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)
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
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
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
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
Modern day
Now thought that protons and neutrons are made up of even smaller particles called quarks
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
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
Isotopes
Different masses
Same no. of protons
Same no. of electrons
Different number of neutrons
Atomic number
Proton number
Mass number
Protons + neutrons
Relative molecular mass, Mr
Add together the relative atomic masses of each atom making up a molecule
Relative formula mass
Add together the relative atomic masses of each atom making up the formula unit
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
A mass spectrum shows
Relative or percentage abundance on the y-axis and mass-to-charge ratio on the x-axis
How is mass-to-charge ratio in all mass spectra shown as
m/z
m= mass z= the charge of the ion (usually 1)
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
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
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
Ammonium
NH4+
Hydroxide
OH-
Nitrate
NO3-
Carbonate
CO3^2-
Sulphate
SO4^2-
Avogadro’s constant
6.02 times 10^23 mol^-1
Number of moles
Mass divided by molar mass
Empirical formula
Write out the atoms present in the compound
Divide the amount of each element present by its molar mass (this gives you the molar ratio)
Divide these values by the smallest value to simplify the ratio
Ensure your answer is the simplest whole number ratio
Write out the empirical formula
At RTP, one mole of gas molecules occupies approximately
24.0 dm^3
24,000 cm^3
n=
v DIVIDED BY 24.0 (dm^3)
For gases
n=
v DIVIDED BY 24,000 (cm^3)
For gases
Gases are assumed to behave in an ideal way:
In continuous motion
Don’t experience any intermolecular forces
Exert pressure when they collide with each other and the container walls
All collisions are elastic - no K.E. lost
K.E. inc. with inc. temp.
Gas molecules are so small so any differences in sizes of different gas molecules can be ignored
Ideal gas equation:
pV = nRT
p= pressure (Pascal, Pa. 1 atm =101325 Pa)
V= volume (Cubic metres, m^3. 1 m^3 = 1000 dm^3)
n= number of moles
R= gas constant (8.314 J mol^-1 K^-1)
T= temperature (Kelvin, K. 0*C = 273 K)
Concentration tells you
How much solute is dissolved in a given amount of solvent
Unit: mol dm^-3
n=
c TIMES BY v (in dm^3)
For solutions
n=
c TIMES BY V DIVIDED BY 1000 (in cm^3)
For solutions
In order to make a standard solution, you will need to..
Know the concentration and volume of the solution you need to make
Work out the amount, in mol, of the solute needed
Convert this amount of solute into a mass, in g, so you know how much to weigh out
Step-by-step of how you make up a standard solution
Step 1: Accurately weigh the required mass of solute. Use a beaker/weigh boat, a balance and a spatula.
Step 2: Transfer the solute into a beaker. Thoroughly wash the weigh boat with distilled water. Add enough distilled water to the beaker to dissolve the solute. Use a glass rod to mix the solution until all the solute dissolves.
Step 3: Use distilled water to wash out the contents of the beaker and pour all the contents into a volumetric flask. Rinse the glass rod with distilled water too and add that to the volumetric flask.
Step 4: Fill up the volumetric flask with distilled water. Stop just before the line on the flask.
Step 5: Use a pipette to get up to the required volume. The bottom of the meniscus needs to touch the line on the flask. If you go over the line then you have to start again!
Step 6: Put the stopper in the flask. Invert the flask multiple times.
Mass concentration are measured in the units
g dm^-3
M refers to a solution with a concentration in
moles per cubic decimetre
Giant structures
Ionic compounds
Giant covalent structures
All metals
Some non-metals e.g. carbon, silicon and boron
Percentage yield
Actual amount in mol of product DIVIDED BY theoretical amount in mol of product
TIMES BY 100
Atom economy
Molecular mass of the desired product DIVIDED BY the sum of the molecular masses of all the products
TIMES BY 100
Sulfuric acid
H2SO4
Hydrochloric acid
HCl
Nitric acid
HNO3
Ethanoic (acetic) acid
CH3COOH
An acid is
A proton donor
Strong acids
V good at giving up H+ ions
Dissociate fully or almost fully
Weak acids
Not v good at giving up H+ ions
Only partially dissociate
Common bases
Metal oxides and hydroxides
E.g. MgO, CuO, NaOH, Mg(OH)2
Ammonia (base)
Ammonia = NH3
Amines = CH3NH2
Magnesium hydroxide, Mg(OH)2 use
“Milk of magnesia”
Used for treating acid indigestion
Calcium hydroxide, Ca(OH)2 use
Also known as lime
Used for reducing the acidity of acid soils
A base
Is a proton, H+ acceptor
Bases neutralise acids
Alkalis
Release OH-(aq) ions when they dissolve in water
Amphoteric substances
Behave as acids and bases
Salts
Ionic compounds
Has a cation, usually a metal or ammonium ion
Has an anion, derived from the acid
The formula of the salt is the same as that of the parent acid, except that the H+ ion has been replaced by the positive ion
Salts can be produced by neutralising acids with bases (proton acceptors) e.g.
Carbonates
Metal oxides
Alkalis
BASHO
Base + acid –> salt + water
MASH
Metal + acid –> salt + hydrogen
CASHOCO
Carbonate + acid –> salt + water + carbon dioxide
Hydrous
The crystalline form containing water
Anhydrous
The crystalline form containing no water
Dot formula
Gives the ratio between the number of compound molecules and the number of water molecules within the crystalline structure
In a titration you MUST know
The concentration of one of your solutions
The unknown info in a titration may be
The conc. of the solution
A molar mass
A formula
The number of molecules of water of crystallisation
How to carry out a titration
- Using a pipette, add a measured volume of one solution to a conical flask. Add a suitable indicator
- Place the other solution in a burette
- Add the solution in the burette to the solution in the conical flask until the reaction has just been completed (end point). Measure the volume of solution added from the burette
- You now know the volume of one solution that exactly reacts with the volume of the other solution
Methyl orange
Red in acid
Yellow in base
Orange is end point colour
Bromothymol blue
Yellow in acid
Blue in base
Green is end point colour
Phenolphthalein
Colourless in acid
Pink in base
Pale pink is end point colour*
*assuming that the aq base has been added from the burette to the aq acid. If the acid is added to the base, the titration is complete when the solution goes colourless.
Oxidation number of uncombined element
0
Oxidation number of oxygen in peroxides
-1
Oxidation number of combined hydrogen
+1
Oxidation number of combined hydrogen in metal hydrides
-1
Oxidation number of simple ions
The charge on the ion
Oxidation number of combined fluorine
-1
Oxidation numbers in compounds
The sum of the oxidation numbers must equal the overall charge of 0
Oxidation numbers in molecular ions
The sum of the oxidation numbers must equal the overall charge
Transition elements form ions with
Different oxidation numbers
Oxyanions
Negative ions that contain an element along with oxygen
Oxidation
Gain of oxygen
Loss of electrons
An increase in oxidation number
Reduction is
Loss of oxygen
Gain of electrons
A decrease in oxidation number
In redox reactions of metals with acids
The metal is oxidised, forming positive metal ions
The hydrogen in the acid is reduced, forming the element hydrogen as a gas (H2)
